Chemistrynotes

Periodic Table of the Elements

Legend
Metals	A solid substance that is a good conductor of heat and electricity. Can be formed into many shapes.
Metalloid	"Middle elements" - conduct heat and electricity better than nonmetals, but not as well as metals. Easier to shape than nonmetals, but not as easy as metals. Solid at room temperature.
Nonmetals	A poor conductor of heat and electricity. Not easily formed into shapes.

This notes section will be updated and improved in the first few weeks of school.

Chapter 2: Matter and Change

What is matter? Anything that has mass and takes up space.
States of matter: Solid, Liquid and Gas
Gas (limited to those substances that exist in the gaseous state at ordinary room temperature)
Vapor (the gaseous state of a substance that is generally
Physical changes: matter can be changed in various ways without altering the chemical composition (i.e.: cutting, grinding, changes in temperature, etc…).
Mixtures-Definition: a physical blend of two or more substances.
Classifying mixtures:
1) Heterogeneous: one that is not uniform in composition
2) Homogeneous: one that is uniform in composition
Solutions: a homogeneous mixture that can be gas, liquid or solid.
Solution Examples
Gas-Carbon dioxide and oxygen in nitrogen (air)
Liquid-Acetic acid in water (vinegar)
Solid-Copper in silver (alloy, silver)
Phase: any part of a system with uniform composition (i.e.: a homogenous mixture consists of one phase while a heterogeneous mixture consists of two or more phases)
Separating Mixtures
Distillation: a liquid is boiled to produce a vapor and is then condensed again to a liquid.
Elements and compounds
Elements: simplest form of matter
Compounds: two or more elements combined
Symbols and Formulas
Each element is represented by one or two letter symbols.
Each symbol represents a short hand way to write chemical formulas of compounds.
o Water = H₂0
o Table sugar = C₁₂H₂₂O₁₁
Chemical Reactions
Changing chemical reactants to products
In chemical reactions the starting substances are called reactants and the substances formed are called products.
In a chemical reaction one or more substances are changed to new substances.
Chemical property: the ability of a substance to undergo a chemical reaction and to form a new substance (i.e.: rusting is a chemical property of iron).
Conservation of mass: massproducts = massreactants

Chapter 3: Scientific Measurement

Section 3.1 - The Importance of Measurement
Qualitative and Quantitative Measurements
- Qualitative measurements give results in a descriptive, nonnumerical form.
- Quantitative measurements give results in a definite form, usually as numbers and units.

Scientific Notation
- Scientific Notation is a number expressed as the product of two numbers: a coefficient and 10 raised to a power (i.e.: 36 000 = 3.6 x 10⁴).

- Multiplication and Division Using Scientific Notation
- Multiplication: multiply the coefficient and add the exponent

Example: (2.0 x 10⁴)x(3.0x10⁶) = 6.0 x 10¹⁰

- Division: divide the coefficient and subtract the exponent

Example: 8.0 x 10⁶ = 4.0 x 10¹
2.0 x 10⁵

- Addition and Subtraction Using Scientific Notation
- Before adding or subtracting you must make the exponents the same because they determine the location of the decimal points

Example: 2.0 x 10⁶ + 6.5 x 10⁵
= 2.0 x 10⁶ + 0.65 x 10⁶
= 1.35 x 10⁶

Section 3.2 - Uncertainty in Measurements
Accuracy, Precision, and Error
- Accuracy is how close a measurement comes to the actual or true value (i.e.: hitting the bullseye every time).

- Precision is how close a series of measurements are to each other (i.e.: obtaining similar results every time).

- Error = accepted value - experimental value

- % Error = (error) divided by the accepted value x 100%

Significant Figures in Measurement
- Significant figures in a measurement include all of the figures that are known, plus the last digit that is estimated.
- It is important to report the correct number of significant figures because calculated answers depend on the number of significant figures.
- Rules
1. Every non-zero digit in a reported measurement is assumed to be significant.

Example: 23, 987 = 5 significant digits

2. Zeros between non-zero measurements are significant.

Example: 304 = 3 significant digits

3. Left most zero digits appearing in front of non-zero figures are not significant; they are place holders.

Example: 0.0038 = 2 significant digits

4. Zeros at the end of a number and to the right of the decimal point are significant.

Example: 9.0 = 2 significant digits

5. Zeros at the right most end of a measurement that lie to the left of a decimal point are not significant; they are place holders.

Example: 7000 = 1 significant digit

6. There are unlimited significant figures when you are counting.
Example: 23 people = 23 people not 22.9

7. There are unlimited significant figures when using exactly defined quantities.

Example: 60 minutes = 1 hour

Significant Figures in Calculations

- Rounding: 7.7 x 5.4 = 41.58
- Round to 42 because the answer can not be more precise than the measured values (i.e.: 2 significant figures).
- Round up if the last number is greater than 5 and down if it is less than 5.

- Addition and Subtraction:
- Numbers should be rounded to the same number of decimal places.

Example: 2.1 - 0.98 = 1.12

- Multiplication and Division:
- Significant figures have nothing to do with the position of the decimal point. Round answer to the least number of significant figures.

Example: 4521.2 = 19.3213675 = 19.3
234

Section 3.3 - International System of Units
Units of Length
- The metric system is important because of its simplicity (i.e.: it is base on multiples of 10's).
- Important SI units include:
Length (m)
Volume (m3)
Mass (kg)
Density (g/cm3)
Temperature (K)
Amount of Substance (mol)
Pressure (Pa)

- Important prefixes include:
Kilo = 1000
deci = 0.1
centi = 0.01
milli = 0.001
micro = 0.000 001

Units of Volume
- Important units include:
Liter (L)
milliliter (mL)
cubic centiliter (cm3)
- Note: 1000 mL = 1 L
1 cm3 = 1 mL

Units of Mass
- Important units include:
Kilogram (kg)
gram (g)
milligram (mg)
- Note: 1 kg = 1000 g
1 g = 1000 mg

Section 3.4 - Density
Determining Density

Density = mass divided by volume
mass is usually in grams and Volume in cubic centimeters
Example: Which has less density 1 lb of lead or
1 lb feathers? Explain.

Specific Gravity

Specific Gravity = density of substance divided by the density of water

- measured with a hydrometer

Section 3.5 - Temperature
Measuring Temperature
- Temperature determines the direction of heat transfer.
- Heat will move from a higher to lower temperature
Example:When your mom screams at you to
close the door because you're letting the
cold in is that statement correct? Explain.

- Temperature scales include degrees Celsius or Centigrade (i.e.: 0° = freezing temperature of water and 100°= boiling temperature of water) and Kelvin (i.e.: K° = C° + 273)

Chapter 4: Problem Solving in Chemistry
What Do I Know?
Problem solving involves developing a plan: A Three-Step Problem-Solving Approach: (1) ANALYZE, (2) CALCULATE, and (3) EVALUATE
Conversion Factors
A ratio of equivalent measurements, such as 1m/100cm or 100cm/1m is called a conversion factor.
Conversion factors are useful in solving problems in which a given measurement must be expressed in some other unit.
Dimensional Analysis
Dimensional Analysis is a way to analyze and solve problems using units, or dimensions, of the measurements (3tsp = 1Tbsp).
Multi-Step Problems
When converting between units it is often necessary to use more than one conversion factor.
Converting Complex Units
Complex units include: densities of solids, liquids (g/cm3), gas mileage (mi/gal) etc…Converting these units is just as easy but it will take more steps.

Content Outline:

Atoms Chapter 5

Early Models of the Atom
- Democritus first suggested the existence of atoms in the 4th century B.C. He believed these atoms were invisible an indestructible.

- Although Democritus’s ideas agreed with later scientific theory they did not explain chemical behaviour.

- Dalton, and English schoolteacher, performed experiments to test and correct his atomic theory. He studied the ratios in which elements combine in chemical reactions. Based on his results he developed Dalton’s Atomic Theory:

1. All elements are composed of tiny invisible particles called atoms.

2. Atoms of the same element are identical. The atoms of any one element are different from those of any other element.

3. Atoms of different elements can be physically mixed together or can chemically combine with one another in simple whole-number ratios to form compounds.

4. Chemical reactions occur when atoms are separated, joined, or rearranged. Atoms of one element, however, are never changed into atoms of another element as a result of a chemical reaction.

Just How Small is and Atom?
- Example: If you grind a copper penny into dust it is still composed of copper. If you continue to make the copper dust into smaller and smaller particles you would eventually come upon a speck copper that could no longer be divided. This speck would be the atom.

- Atom: The smallest particle of an element that retains the properties of that element.

- A scanning tunneling microscope makes seeing individual atoms possible.

Structure of the Nuclear Atom
Electrons
- Most of Dalton’s theory is accepted today. The one change is that atoms are now known to be divisible (i.e.: they can be broken into smaller, more fundamental particles).

- Electrons: are negatively charged subatomic particles.

- J.J. Thomson discovered electrons in 1897 by passing an electric current through a gas at low pressure. The gases were sealed in glass tubes fitted with metal disks at both ends. One electrode, the anode, became negatively charged while the other, the cathode, became positively charged. A glowing beam (the cathode ray) flowed between the electrodes.

- Thompson discovered that electrons were negatively charged because a negative electrical ray repelled the ray while a positive electrical charge attracted it.

Protons and Neutrons
- If cathode rays are electrons given off by atoms then what remains?

- Atoms have no net electric charge

- Charges are carried by particles of matter.

- Electric charges always exist in whole-number multiples

- When an equal number of negatively charged particles combine with an equal number of positively charged particles and electrically neutral charged particle remains.

- Therefore, a particle unit of positive charge should remain when an atom looses and electron.

- Goldstein discovered the proton (a positively charged particle) when he noticed that rays were travelling in the reverse direction of the cathode ray.

- Chadwick determined the existence of neutrons. The neutron is a subatomic particle with no charge but with a mass nearly equal to that of the proton.

The Atomic Nucleus
- Rutherford performed an experiment to test how these subatomic particles were put together in the atom. He shot alpha particles (i.e.: helium with no electrons) at a thin sheet of gold foil. According to Dalton’s atomic theory, Rutherford surmised that the alpha particle would pass easily through the thin sheet with only a slight deflection. It was found that the alpha particles passed directly through the foil with no deflection and a small fraction of the alpha particles bounced back at very large angles. Based on these results Rutherford suggested a new theory of the atom in which he proposed that the atom is mostly empty space. He proposed that the positive charge and most of the mass were concentrated in a small region that has enough positive charge to account for the great deflection of some of the alpha particles. He called this region the nucleus.

- Nucleus: The central core of the atom comprised of protons and neutrons.

Distinguishing Between Atoms

Atomic Number
- Atoms are made up of electrons, protons, and neutrons.

- Protons and neutrons make up the dense nucleus while electrons surround the nucleus and account for most of the atom.

- Elements will differ because they contain different numbers of protons (i.e.: Hydrogen had one proton while oxygen has eight).

- The atomic number of an element is the number of protons in the nucleus. Note that the number of protons will equal the number of electrons.

Mass Number
- Mass Number: The total number of protons and neutrons in an atom.

- You can determine the composition of the atom if you know the atomic number and the mass number of any atom (i.e.: atomic number is equal to the number of protons, which equals the number of electrons. The mass number is equal to the number of protons plus the number of neutrons.

Isotopes
- Isotopes: atoms that have the same number of protons but different numbers of neutrons.

- Because isotopes of an element have different numbers of neutrons they also have different mass numbers.

- Despite these differences, isotopes are chemically alike since the number of protons and electrons have not changed.

- Example: three types of hydrogen include the common form of hydrogen, hydrogen-2 or deuterium with one neutron and hydrogen-3 or tritium with two neutrons.

Atomic Mass
- Mass of a proton or a neutron is 1.67 x 10-24 g.

- Mass of an electron is 9.11 x 10-28g.

- Since the 1920’s we’ve been able to measure these tiny masses by using a mass spectrometer.

- Since these masses are so small it is more convenient to compare the relative masses of atoms using a reference isotope. The isotope chosen is carbon-12.

- Carbon-12 was designated an atomic mass unit (amu) of 12.

- The atomic mass unit is defined as one-twelfth the mass of a carbon-12 atom.

- Helium-4 with a mass of 4.0026 amu has a mass of one-third the mass of carbon-12. Nickel-60 has an atomic mass of 5 times that of carbon-12.

- Atomic mass is a weighted average mass of the atoms in a naturally occurring sample of the element. A weighted average reflects both the mass and the relative abundance of the isotopes as they occur in nature.

- To calculate atomic mass of an element as a weighted average of the masses of its isotopes you need to know:

1. The number of stable isotopes of the element.

2. The mass of each isotope.

3. The natural percent abundance of each isotope.

The Periodic Table
Development of the Periodic Table
- About 70 elements had been discovered by the mid-1800’s but no one had found a way to relate the elements in a systematic way until Dmitri Mendeleev.

- He listed the elements in columns in order of increasing atomic mass. He then arranged the columns so that the elements with the most similar properties were side by side.

- In 1913, Henry Moseley determined the atomic number of the atoms of the elements. He arranged the elements in a table by order of atomic number instead of atomic mass. This is the way the periodic table is arranged today.

The Modern Periodic Table
- Each element is identified by its symbol placed in a square.

- The atomic number of the element is shown centered above the symbol.

- The atomic mass and the name of the element are shown below the symbol.

- The elements are listed in order of increasing atomic number from left to right and from top to bottom.

- The horizontal rows of the periodic table are known as Periods (there are seven in total).

- Moving from one period to the next gives rise to the periodic law. When the elements are arranged in order of increasing atomic number there is a periodic repetition of their physical properties and chemical properties. Elements with similar physical and chemical properties end up in the same column.

- Every vertical column is known as a group or a family. A number and a letter A or B identify each group (i.e.: the first column is labeled group 1A. All elements in this group excluding hydrogen react vigorously with water.)

- Group A elements are made up of Group 1A through 7A and Group 0. Group A elements are called representative elements because they exhibit a wide range of both physical and chemical properties.

- Representative elements can be divided into metals, which have high electric conductivity and high luster. Except for hydrogen the elements on the left of the periodic table are metals. Groups 1A are known as alkali metals and Group 2A are known as alkaline earth metals.

- Most of the remaining elements are not of Group A. These include transition metals and inner transition metals. Together these make up Group B.

- The non-metals occupy the upper right corner of the periodic table. Nonmetals are elements that are generally nonlustrous and are generally poor conductors of electricity.

- The nonmetals of Group 7A are called halogens. Group 0 elements are known as noble gases, which are also referred to as inert gases because they undergo few chemical reactions.

- Notice that the heavier stair step line divides the metals from the nonmetals. Most of the elements that border this line are called metalloids. These metals have properties that are intermediate between metals and nonmetals.

Content Outline:

Chapter 13

Models of the atom:

The Evolution of Atomic Models:

- The atomic model thus far has been assuming a combination of protons and neutrons making up the nucleus and the electrons surrounding the nucleus.

PROBLEM:

o Does not explain why metals or components of metals give off characteristics colors when heated in a flame.

- Chemical properties of atoms, ions, and molecules are related to the arrangement of the electrons. This we further expand upon.

Dalton Atomic Theory: Believed the atom was a solid invisible mass

- The discovery of subatomic particles dismissed every theory scientists had about invisible mass.

J.J. Thompson: proposed a revised model of the atom based on his discovery of the electron (referred to as the plum-pudding model).

Plum-pudding model:

- Describes the atom as having negatively charged electrons stuck on a positively charged lump.

PROBLEMS:

o It did not take into account the number of protons and electrons, their arrangements in the atom, or the ease with which the atoms are stripped of its electrons.

Ernest Rutherford: based on Thompson’s discoveries he proposed the nuclear atom in which electrons surrounded a dense atom.

- Later experiments proposed that the nuclei of the atom was made up of protons ( + charge) and neutrons ( - charge).

Neils Bohr: proposed that the electrons are arranged in concentric circular path (i.e. orbital) around the nucleus. Also known as planetary model.

- Proposed that electrons in a path have a fixed energy, they do not loose energy and they can not fall into the nucleus.

- Energy level: the region around the nucleus where the electron in known to be moving.

o The lowest energy level (1s) corresponds to the lowest energy. The highest (7p) corresponds to the most energy.

o Electrons can jump from one energy field to another, however, IT CANNOT ORBIT BETWEEN LEVELS.

o In order to move from one energy level to the next, electrons must either gain or lose the right amount of energy. This energy is known as Quantum Energy.

n = 4

n =5

Energy Levels:

n = 3

Increasing

n = 2

Energy Level

n = 1

The Quantum Mechanical Model:

- Previous models have been physical models based on the motion of large objects. The quantum model is primarily mathematical.

- Like the Bohr model, the quantum model restricts the energy of electrons to certain values. However, the quantum model does not define the exact travel path of the electrons (i.e. estimates the probability of finding an electron).

EXAMPLE:

Probability: If you place three red marbles and one green marble into a box, and then pick up one marble, the probability of picking a green one is 1 in 4 or 25%.

- The space around the nucleus is assumed to be a cloud; the higher the probability of finding an electron the thicker (or more dense) the cloud.

Atomic Orbitals:

- The quantum mechanical model designates energy levels of electrons by principal quantum numbers (n) (i.e. n – 1,2,3,4. etc…).

- Distances from the nucleus increases as n increases.

- Within each energy level there are sublevels (Regions in which the electrons is likely to be found).

- Letters denote atomic orbitals.

o s orbitals - sphere

o p orbitals – dumbbell shaped

o d orbitals – 4 of the 5 have clover leaf shapes

o f orbitals are hard to visualize

o p and d orbitals are close to the nucleus where the probability of finding an electron is law.

Electron Configuration: the way in which the electrons are arranged around the nucleus

- In most natural phenomena, change proceeds toward the lowest possible energy. High-energy levels are unstable.

- Unstable systems will lose energy in order to become more stable.

- In the atom, electrons and the nucleus will interact in order to make the most stable arrangements possible.

Aufbau Principle: electrons enter the lowest energy orbitals first

- The various orbitals within a sublevel of a principal energy level are always of equal energy

- Within an energy level, the sublevel s is always the lowest energy level

- Energy levels can overlap adjacent energy levels

- Electrons enter the lowest energy first

Pauli Exclusion Principle: an atomic orbital may describe at most two electrons

- to occupy the same orbital, the two electrons must have opposite spins

- the orbital containing paired electrons is written

Hund’s rule: when electrons occupy orbital of equal energy, one electron enters each orbital until all the orbitals contain two electrons with two spins

Denoting Electron Configurations: there are two ways

1) 1) Uses opposing arrows to denote electrons

2) A short hand formula (i.e.: 1s22s2)

Exceptional Electron Configuration:

- You can obtain correct electron configuration for the elements up to vanadium by following the Aufbau diagram for orbital filling.

Physics and the Quantum Mechanical Model:

Light and atomic Spectra:

- The quantum mechanical model of the atom grew out of the study of light

- According to the wave model light consists of electromagnetic waves

Electromagnetic Radiation: includes radio waves, microwaves, infrared waves, visible light, ultraviolet waves, x-rays, and gamma rays

- electromagnetic waves travel at 3.0x108 m/s or 3.0 x 1010 cm/s(speed of light)

Amplitude: the waves height from origin to crest

Wavelength: (l) the distance between crests

Frequency: (u) the number of wave crests to pass a given point over unit of time (units given in Hz <hertz> or s –1)

c = lu

- the larger the wavelength the higher the frequency

Spectrum: a rainbow is an example of a spectrum

- Every element will emit light when it is excited by the passage of an electric discharge through its gas or vapor

Atomic Emission Spectrum: occurs by passing a light emitted by an element through a

prism

- When a gaseous or vaporized element is excited by electricity, the spectrum of the emitted light should be continuous

- Classical physics does not explain the emission spectra of atoms which consist of lines

- Planck showed mathematically that the amount of radiation energy (E) absorbed or emitted by a body is proportionally to the frequency of the radiation

H = Planck’s constant = 6.6262x10-34 Jxs

- Planck proposed that quanta of energy are absorbed or emitted

- Einstein proposed that light could be described as quanta of energy that behaved as particles

- Light quanta are called photons

Photoelectric Effect: metals eject electrons called photoelectrons when light shines on

them. Alkali metals are particularly subject to it but not all light

causes this effect

Atomic Spectra:

- exciting an electron will cause it to jump to the next energy level. Energy will be absorbed.

- Only electrons that move from a higher to a lower energy level will emit light because they are losing energy

Lyman Series: electrons move from higher energy level to n=1

Balmer Series: electrons move from higher energy levels to n=2

Paschen Series: electrons move from higher energy levels to n=3

Quantum Mechanics:

- energy is absorbed and emitted in packages

- light behaves as both a wave and particle

· In 1924, Louis DeBroglie, derived an equation that described the wavelength of moving particles:

l = h

where h is Planck's constant, m is the mass of the particle and v is the velocity of the particle.

· DeBroglie's equation predicts that all matter exhibits wave like and particle like properties.

· Classical vs. Quantum mechanics:

Classical mechanic adequately describes the motion of bodies much larger than the atoms they are comprised of. It appears that such a body gains or loses energy.
Quantum mechanics describes the motion of subatomic particles and atoms as waves. These particles gain or lose energy in packets called quanta.
· Heisenberg Uncertainty Principle: it is impossible to know exactly both the velocity and the position of a particle at the same time. This principle is much more obvious with smaller particles than with large ones.

Content Outline:

Chapter 14

Classifying Elements by Chemical Configuration

1) Electrons play a significant role in determining physical and chemical properties of elements.

2) Arrangements of elements on the periodic table depend on such properties.

3) Elements are classified into four categories based on their electron configuration.

1) The Noble Gases

- Elements in which the outer most s and p sublevels are filled (Group 0).

i.e.: Helium 1s2

Neon 1s22s22p6

- These gases DO NOT participate in chemical reactions!!!

2) The Representative Elements

- In these elements the outermost s and p sublevels are only partially filled. Groups of representative elements include Group 1A, Group 2A and Group 7A.

- *For any representative element the group number equals the number of electrons in the outermost energy level.

i.e.: Group 1A elements have “1” electron in the outermost shell.

Lithium 1s22s2

Sodium 1s22s22p63s1

Group 4A have “4” electrons in the outermost shell.

Carbon 1s22s22p2

Silicon 1s22s22p63s23p2

3) The Transition Metals

- The outermost s sublevel and nearby d sublevel contain electrons. These Group B elements are characterized by addition of electrons to the d orbitals.

4) The Inner Transition Metals

- In these metallic elements the outermost s sublevel and nearby f sublevel generally contain electrons.

- The periodic table can be divided into sections or blocks. The s block contains s1 and s2 configurations and is made up of Group 1A and 2A elements. The p block is composed of Group 3A, 4A, 5A, 6A, 7A and 0 elements (DOES NOT include Helium).

- In order to determine electron configurations you simply need to read the periodic table from left to right and from top to bottom (just like a book).

1) Each period corresponds to the principle energy level.

2) The number of electrons in each sublevel is determined by counting over to the element (moving from left to right).

3) For transition elements, electrons are added to a d sublevel with a principle energy level that is one less than the period number.

4) For inner transition metals, the principle energy level of the f sublevel is two less than the period number.

Trends in Atomic Size

- In chapter 13 we learned that atoms do not have a defined limit of its size, therefore, we can not measure the atoms radius.

- Methods involved in “estimating” the size of an atom include:

1) Solid Crystalline Structured Atoms: using x-ray diffraction can provide an estimate of the distance between the nuclei making up the crystalline structure.

2) Diatomic Models: the atomic radius is one-half the distance between the nuclei in the diatomic molecules.

The Periodic Table

1) Group Trends: atomic size generally increases as you move down a group (i.e.: electrons will add to successive higher principle energy level and nuclear charge increases, as a result, shielding the nucleus)/

2) Periodic Trends: atomic size generally decreases as you move from left to right (i.e.: principle energy levels will remain the same. The effect of the increased nuclear charge on the outer most electrons is to pull them closer to the nucleus).

Trends in Ionization Energy

- * When an atom gains or loses an electron it becomes an ion.

- The energy required to overcome the attraction of nuclear charge and remove an electron from a gaseous atom is called the ionization energy.

- * Removing an electron results in a 1+ charge.

- * Adding an electron results in a 1- charge.

- The energy required to remove the first outermost electron is called the first ionization energy. Removing the outermost electron from the gaseous 1+ ion requires a second ionization energy and so on.

1) Group Trends: the first ionization energy generally decreases as you move down the group (i.e.: since atomic size increases as you move down a group the outermost electrons are further from the nucleus and, therefore, easier to remove).

2) Periodic Trends: for representative elements, the first ionization energy generally increases as you move from left to right (i.e.: since nuclear charge increases and the shielding effect is constant as you move across an increased attraction between the nucleus an electron results).

Trends in Ion Size

- Atoms of metallic elements have low ionization energies (i.e.: they form positive ions easily).

- Nonmetallic elements readily form negative ions.

1) Group Trends: Positive ions are always smaller than the neutral atoms (i.e.: losing outershell electrons increases the attraction between the nucleus and the remaining electrons).

2) Periodic Trends: Going from left to right the size of positive ions decreases.

Trends in Electronegativity

- Electronegativity: the tendency for the atoms of the element to attract electrons when they are chemically combined with atoms of another element.

- Moving from left to right, electronegativity will increase (Elements at the far left of the table, therefore, have low electronegativity [excluding noble gases] and elements to the far right have high electronegativity).

Chapter 15: Ionic Bonding and Ionic Compounds

Electron Configuration in Ionic Bonding

Valence Electrons
- Valence Electrons: the electrons in the highest occupied energy level of an element’s atom.

- Valence electrons determine why atoms combine to form specific compounds and why their formulas are what they are.

- Elements within each group in the periodic table will behave similarly because they have the same number of valence electrons.

- The number of valence electron will determine an elements chemical and physical properties.

- Determining the number of valence electrons can be done by determining the atoms electron configuration.

- The number of valence electrons are related to the group numbers in the periodic table (i.e.: in order to determine the number of valence electrons for a particular element look at its group number – Hydrogen, Lithium, Sodium, etc.. will all have one valence electron). Group 0 elements will be the only exception to this rule. Helium has 2 valence electrons while all other noble gases have 8.

- Valence electrons will normally be the only electrons used in chemical bonding.

- Electron Dot Structures: diagrams that show valence electrons as dots. The inner electrons and the atomic nuclei are included in the letter symbol for the element being represented.

Electron Configurations for Cations
- Octet Rule: in forming compounds, atoms tend to achieve the electron configuration of a noble gas. An octet is a set of eight (i.e.: Each noble gas had 8 valence electrons).

- Metallic atoms will tend to lose their valence electrons therefore leaving a complete octet in the next lowest energy level.

- Atoms of some nonmetallic atoms will tend to gain electrons of share electrons with another nonmetallic atoms in order to complete an octet.

- The octet rule applies to most atoms in compounds.

- Cations: are formed when an atom loses a valence electron. The lose of an electron will result in a positively charge ion.

- Cations in Group 1A will always have a 1+ charge because they have 1 valence electron. Those in Group 2A will always have a 2+ charge because they have 2 valence electrons.

- Transition metals tend to have various charges (i.e.: Fe(II) or (III)) this is because some transition metals electron configurations do not have noble gas electron configurations and as a result are an exception to the octet rule. In order for silver to achieve the structure of krypton it would have to loose 11 electrons. In order for silver to achieve the configuration of xenon it would have to gain 7 electron. Chances of this happening are unlikely therefore silver can not reach noble gas configuration. Silver would then lose 1 electron (5s1) in order to reach the next favorable configuration.

Electron Configurations for Anions
- Anion: an atom or a group of atoms with a negative charge. Anions are formed when an atom acquires a negatively charged electron.

- Atoms of nonmetallic elements attain noble gas electron configuration more easily by gaining electrons than by losing then so nonmetallic elements will tend to anions. By gaining electrons these nonmetallic elements will then have a full octet in their outer shell.

- Halide Ions: produced when atoms of chlorine and other halogens gain electrons. All halide ions have 7 valence electrons and need to obtain one electron to achieve noble gas electron configuration.

Ionic Bonds
Formation of Ionic Compounds
- Anions and cations, being of opposite charge, will attract each other by electrostatic forces.

- Ionic Bonds: the forces of attraction that bind these oppositely charge ions.

- Ionic Compounds: compound that consist of electrically neutral groups of ions joined by electrostatic forces. In an ionic compound the total positive charges must equal the total negative charges.

- Example:

- Sodium has one valence electron that it can afford to lose in order to reach noble gas electron configuration. Chlorine had 7 valence electrons therefore it can afford to gain one electron in order to reach noble gas electron configuration. When sodium and chlorine react to form a compound sodium therefore give chlorine its extra electron.

Properties of Ionic Compounds
- At room temperature most ionic compounds are crystalline solids. The ions will therefore be arranged in repetitive 3D patterns.

- Coordinate Number: the number of ions of opposite charge that surround the ion in a crystal.

- Crystal structure is determined using x-ray diffraction. X-rays will pass through a crystal and the pattern will be exposed on film as a result of the x-rays being deflected by the crystal structure.

Bonding in Metals
Metallic Bonds and Metallic Properties
- Metals are made up of closely packed cations as opposed to neutral atoms. The cations are surrounded by mobile valence electrons.

- Metallic Bonds: consist of the attraction of the free-floating valence electrons for the positively charged metal ions. These bonds are the forces of attraction that hold metals together.

- Metals are good conductors of electricity because electrons can flow freely in them.

Crystalline Structure of Metals
- In a body centered cubic structure every atom has eight neighbors.

- In a face centered cubic structure every atom has twelve neighbors.

- In a hexagonal closed packet arrangements also have twelve neighbors however they are not limited strictly to metal atoms.

Alloys
- Alloys are mixtures of two or more elements, at least one of which is a metal.

- Alloys are important because their properties are often superior to those of their component elements (i.e.: Sterling silver is stronger than pure silver). The most important alloy in use today is steel.

- Alloys can form from their component atoms in various ways.

- Substitution Alloy: If the atoms of the components in an alloy are about the same size they can replace each other.

Interstitial Alloy: if the atomic sizes are quite different the smaller atoms can fit into the spaces between the larger atoms.

Test #2 Study Guide

Chapter 5

KEY TERMS:

- Alkali metals - Cathode ray - Isotope Nonmetal

- Alkaline earth metals - Dalton’s atomic theory - Mass number Nucleus

- Atom - Electron - Metal Period

- Atomic mass - Group - Metalloid Periodic law

- Atomic mass unit - Halogen - Neutron Periodic table

- Atomic number - Inner transition metal - Noble gas Proton

- Representative element

- Transition metal

KEY EQUATIONS AND RELATIONSHIPS
· ATOMIC NUMBER = NUMBER OF PROTONS = NUMBER OF ELECTRONS

· NUMBER OF NEUTRONS = MASS NUMBER – ATOMIC NUMBER

· MASS NUMBER = TOTAL PROTONS + TOTAL NEUTRONS

· ISOTOPES = MASS NUMBER

ATOMIC NUMBER

· ATOMIC MASS UNIT = 1/12 THE MASS OF A CARBON-12 ATOM

· ATOMIC MASS = A WEIGHTED AVERAGE OF AN ATOM IN A NATURALLY

OCCURING SAMPLE OF ELEMENT. *REFLECTS BOTH RELATIVE

· ABUNDANCE AND MASS!!!

LOOK OVER

· Dalton’s atomic theory.

· Describe the size of an atom.

· The difference between protons, electrons, and neutrons in terms of relative mass and charge (i.e.: PROTONS HAVE A POSITIVE CHARGE, NEUTRONS ARE NEUTRAL, AND ELECTRONS HAVE A NEGATIVE CHARGE. mPROTON & NEUTRON=1.67X10-24g, mELECTRON=9.11X10-28g).

· The structure of an atom, including the location of the protons, and neutrons with respect to the nucleus.

· How the atomic number identifies an element (i.e.: ATOMIC NUMBER = NUMBER OF PROTONS = NUMBER OF ELECTRONS).

· How the atomic number and mass number of an element can be used to find the number of protons, electrons and neutrons (i.e.: MASS NUMBER = TOTAL PROTONS + TOTAL NEUTRONS).

· How isotopes differ (i.e.: ATOMS WITH SAME NUMBER OF PROTONS BUT DIFFERENT NUMBER OF NEUTRONS) and why the atomic masses of elements are not whole numbers (i.e.: ATOMIC MASS OF AN ELEMENT IS A WEIGHTED AVERAGE MASS OF THE ATOM IN A NATURALLY OCCURING SAMPLE OF ELEMENT).

· How to calculate the average atomic mass of an element from isotope data (i.e.: NEED TO KNOW THE NUMBER OF STABLE ISOTOPES OF THE ELEMENT, THE MASS OF EACH ISOTOPE, AND THE NATURAL PERCENT ABUNDANCE OF EACH ISOTOPE [LOOK AT THE VEGIUM LAB AND THE ISOTOPE WORKSHEET WE DID]).

· The origin of the periodic table.

· The position of groups (COLUMNS), periods (ROWS), and the transition metals (GROUP B) in the periodic table.

CHAPTER 13

KEY TERMS

- Amplitude - Electron configuration - Hund’s rule - Spectrum

- Atomic Emission - Spectrum Energy level - Pauli exclusion principle - Wavelength (l)

- Atomic Orbital - Frequency Photoelectric effect - Photon

- Aufbau Principle - Ground state - Planck’s constant

- De Broglie’s equation - Heisenberg uncertainty principle - Quantum

- Electromagnetic radiation - Hertz (Hz) - Quantum mechanical model

KEY EQUATIONS
c = lu

E = hv

l = h

LOOK OVER

· The development of the atomic theory.

· The significance of quantized energies of electrons as they relate to the quantum mechanical model (i.e.: THE AMOUNT OF ENERGY REQUIRED TO MOVE AN ELECTRON FROM ITS PRESENT ENERGY LEVEL TO THE NEXT HIGHER ONE).

· How to apply the Aufbau principle(i.e.: ELECTRONS ENTER ORBITALS OF LOWEST ENERGY FIRST), the Pauli exclusion principle (i.e.: AN ATOMIC ORBITAL MAY DESCRIBE AT MOST TWO ELECTRONS), and Hund’s rule (i.e.: WHEN ELECTRONS OCCUPY ORBITALS OF EQUAL ENERGY, ONE ELECTRON ENTERS EACH ORBITAL UNTIL ALL THE ORBITALS CONTAIN ONE ELECTRON WITH PARALLEL SPINS) in writing the electron configuration of elements.

· Why electron configurations for some elements differ from those assigned using the Aufbau principle(i.e.: REMEMBER COPPER AND CHROMIUM ARE SPECIAL CASES – A HALF FULL OR COMPLETELY FULL d LEVEL IS MORE STABLE THAN A PARTIALLY FULL d LEVEL).

· How to calculate the wavelength, frequency, or energy of light, given two of these values (i.e.: YOU WILL BE RESPONSIBLE FOR KNOWING THE THREE KEY EQUATIONS AND BEING ABLE TO SOLVE FOR THE UNKNOWN WHEN TWO VALUES ARE GIVEN. REMEMBER THAT c AND h ARE CONSTANTS [LOOK OVER THE EXAMPLES WE DID IN CLASS THEY WILL HELP YOU).

· The origin of the atomic emission spectrum of an element.

CHAPTER 14

KEY TERMS

- Atomic radius - Noble gases

- Electronegativity - Representative elements

- Inner transition meta - Transition metal

- Ionization energy

LOOK OVER
· The properties of an element based on those of other elements in the periodic table.

· How to use electron configuration to classify elements as noble gases, representative elements, transition metals, or inner transition metals (i.e.: REFER TO FIGURE 14.5. IT SUMMARIZES WHAT YOU SHOULD LOOK FOR WHEN DISTINGUISHING BEWTEEN NOBLE GASES, ETC… ON THE PERIODIC TABLE WITH RESPECT TO ELECTRON CONFIGURATION).

· How to interpret group trends in atomic radii, ionic radii, ionization energies, and electronegativities (i.e.: LOOK AT FIGURE 14.16 IN YOUR TEXTBOOK IT NICELY SUMMARIZES ALL OF THE TRENDS).

CHAPTER 15
KEY TERMS
- Coordination number - Metallic bond

- Electron dot structure - Octet rule

- Halide ion - Valence electron

- Ionic bond

LOOK OVER

· How to use the periodic table to infer the number of valence electrons in an atom and draw it electron dot structure (i.e.: GROUP 1A HAS 1 VALENCE ELECTRON, GROUP 2A HAS 2, GROUP 3A HAS 3, GROUP 4A HAS 4, GROUP 5A HAS 5, GROUP 6A HAS 6, GROUP 7A HAS 7 AND GROUP 0 HAS 8. YOU CAN NOT DETERMINE VALENCE ELECTRONS OF TRANSITION METALS THIS WAY. FOR TRANSITION METALS YOU NEED TO WRITE THE ELECTRON CONFIGURATION. ONLY S AND P SUBLEVELS CAN BE VALENCE ELECTRONS. LOOK AT THE FIGURE YOU WERE GIVEN DURING THE GROUP ACTIVITY ON HOW TO DETERMINE DOT NOTATION. YOU WILL NOT NEED TO DO THE DOT NOTATION FOR TRANSITION METALS).

· The formation of cations from metals (i.e.: METALS WILL MORE READILY GIVE AWAY ELECTRONS) and of anions from nonmetals (i.e.: NONMETALS WILL READILY ACCEPT ELECTRONS).

· The characteristics of an ionic bond (i.e.: ATTRACT ONE ANOTHER BY ELECTROSTATIC FORCES, FORMED BY A METAL AND A NONMETAL, WILL FORM CRYSTAL STRUCTURES).

· The arrangements of atoms in some common metallic crystal structures

Chapter 6 Introduction to Chemical Bonding

Molecules and Molecular Compounds

In nature, noble gases are the only elements that tend to exist in isolation.

Noble gases are also known as monatomic since they consist of single atoms.
A molecule is the smallest electrically neutral unit of a substance that still has the properties of the substance. They are made up of two or more atoms that act as a unit.
Atoms of different elements may combine chemically to form compounds.
Compound Molecules: compounds composed of molecules. Made up of atoms of two or more nonmetals (i.e CO or Carbon monoxide).
Molecular compounds tend to have relatively low melting and boiling points. Many of these compounds therefore exist as gases or liquids at room temperature.

Ions and Ionic Compounds:

Not all compounds are molecular. Many compounds are made up of ions.
Ions: atoms or groups of atoms that have a positive or negative charge. Ions will form when an atom or groups of atoms lose or gain electrons.
Atoms of metallic elements tend to form ions by losing one or more electrons. These are known as cations. Cations will always have a positive charge.
Note that there will be many chemical differences between metals and there cations (i.e.: sodium metal reacts explosively with water but sodium cations are unreactive in water).
Nonmetallic elements tend to form ions by gaining one or more electrons. These are known as anions and have a negative charge.
An anion has more electrons than the electrically neutral atom from which it forms.
Compound composed of cations and anions are known as Ionic Compounds (usually made up of metallic cations and nonmetallic anions).
Ionic compounds are usually solid crystals at room temperature and they melt at high temperatures.

Representing Chemical Compounds

Chemical Formulas:

Chemical Formula:shows the kinds and numbers of atoms in the smallest representative unit of the substance. You can represent the chemical formula of monatomic elements by means of their atomic symbols. If the molecules of the element each have more than one atom a number is then used as a subscript.

Molecular Formulas:

A molecular formula shows the kinds and numbers of atoms present in a molecule of a compound. If there is only one atom, the subscript 1 is omitted.
Note that although a molecular formula shows the composition of a molecule it tells you nothing about its structure.

Formula Units:

Chemical formulas can also be written for ionic compounds. In this case, however, the formula does not represent a molecule.
To represent an ionic compound chemists use a formula unit. This is the lowest whole number ration of ions in the compound. For example, the lowest whole number ratio for NaCl is 1:1.

The Laws of Definite and Multiple Proportions:

Law of Definite Proportions: in samples of any chemical compound, the masses of the elements are always in the same proportions. This law is consistent with Dalton’s Atomic Theory. If atoms combine in simple whole ratios, as Dalton postulates, it follows that their proportions by mass must always be the same.
Law of Multiple Proportions: whenever two elements form more than one compound, the different masses of one element that combine with the same mass of the other element are in the ratio of small whole numbers.

Ionic Charges

Monatomic Ions:

Monatomic Ions:are ions consisting of only one atom, the ionc charges can often be determined by using the periodic table (i.e.: Group 1A elements have a 1+ charge).
Group 4A and Group 0 elements will tend not to form ions.
Many of the cations of transition metals will tend to have multiple charges. There are two methods of naming such cations, however, the preferred method is the Stock System. As part of the name of the element, a Roman numeral in parentheses indicated the numerical value of the charge.
Silver, Cadmium and Zinc will tend not to need roman numberal notation as their charges are normally consistent.

Polyatomic Ions:

Polyatomic Ions: are composed of tightly bound groups of atoms behaving as a unit and carrying a charge.
The names of most polyatomic anions end in -ite or -ate. The exceptions include ammonium, cyanide and hydroxide.
The –ite ending will indicate one less oxygen atom than the –ate ending. Note that the ending will not, however, tell you the actual number of oxygen atoms.

Ionic Compounds

Writing Formulas for Binary Ionic Compounds:

Binary Compounds: compounds made up of two elements.
The positive charge must balance the negative charge when elements form binary compounds (i.e.: the net ionic charge must be zero).
One approach to writing balanced chemical formulas is the crisscross method. In this method the ionic charge of each ion is crossed over and used as a subscript for the other ion.
When the charges are the same (i.e.: Ca²⁺ and S^2-) remember to implement the lowest whole-number ratio. The subscripts would therefore be 1 instead of 2.

Ternary Ionic Compounds:

Ternary Compounds: compounds that contain atoms of three different elements.

Ternary compounds will commonly contain a polyatomic ion.

When naming ternary compounds, the ending given to the polyatomic ion (-ite or -ate) will remain the same. For example, nitrate is a polyatomic ion when it is combined with calcium the name becomes calcium nitrate. The formula for calcium nitrate is written as Ca(NO₃)₂. Notice that the NO₃ is in parenthesis indicating that we need 2 NO₃^- ions in order to balance the charges.

Binary Molecular Compounds:

Unlike binary ionic compounds, binary molecular compounds are made up of two nonmetallic ions.
Prefixes will help in distinguishing between different compounds as follows:

Prefix	Number
mono	1
di-	2
tri-	3
tetra-	4
penta-	5
hexa-	6
hepta-	7
octa-	8
nona-	9
deca-	10

Prefixes will tell the reader how many atoms of each element are present in each molecule of a compound.
Examples: CO = carbon monoxide

CO₂= carbon dioxide

Notice that the second ending will end with –ide just as they do with binary ionic compounds.

Naming Common Acids:

Acids are compounds that produce hydrogen ion when dissolved in water. Some acids include hydrochloric acid, sulfuric acid, and nitric acid.

Summary of Naming and Formula Writing

Rules to Live by:

In an ionic compound, the net ionic charge must equal zero.
An –ide ending generally indicates a binary compound.
An –ite or –ate ending means that there is a polyatomic ion that includes oxygen in the formula.
Prefixes in the name geerally indicate that the compound is molecular. They show the number of each kind of atom in a molecule.

A Roman numberal after the name of a cation shows the ionic charge of the cation.

Chapter 7: Chemical Quantities

What is a mole?

SI measurement of the amount of something (i.e.: particle, mass, volume).

Number of particles in 1 mole?

· 6.02 x 10²³ particles of a substance (Avogadro’s number).

· When you hear someone talk about a representative particle they are normally referring to what is present (i.e.: atoms, molecules, or formula units [ions]).

· Diatomic particles are elements that exist as pairs (i.e.: H₂, N₂, O₂, F₂, Cl₂, Br₂, and Cl₂).

o 1 mole H₂ = 6.02 x 10²³ particles

o 1mole H₂0 = 6.02 x 10²³ particles

o 1 mole Ca²⁺ = 6.02 x 10²³ particles

1. Example:

How many moles of magnesium is 1.23 x 10²³atoms of Mg?

Known:

# of atoms Mg = 1.25 x 10²³atoms Mg

1 mole Mg = 6.02 x 10²³ atoms Mg

# moles Mg = ? mol Mg

1.25 x 10²³atoms Mg x 1 mol Mg = 2.08 x 10^-1 mol Mg

6.02 x 10²³ atoms Mg

2. Example:

How many moles is 2.80 x 10²⁴ atoms of silicon?

Known:

# of atoms Si = 2.80 x 10²⁴ atoms Si

1 mole Si = 6.02 x 10²³ atoms Si

# moles Si = ? mol Si

2.80 x 10²⁴ atoms Si x 1 mol Si = 4.65 mol Si

6.02 x 10²³ atoms Si

3. Example:

How many molecules is 0.360 moles of water?

Known:

# mole H₂O = 0.360 mol H₂O

# molecules H₂O = ? molecules H₂O

0.360 mol H₂O x 6.02 x 10²³ atoms H₂O = 2.167 x 10²³ molecules H₂O

1 mol H₂O

4. Example:

How many atoms is 2.12 moles of C₃H₈?

Known:

1 mole C₃H₈= 6.02 x 10²³ atoms C₃H₈

1 molecule C₃H₈= 11 atoms

2.12 mol C₃H₈x 6.02 x 10²³ atoms C₃H₈x 11 atoms C₃H₈ = 1.40 x 10²⁵ atoms

1mol C₃H₈ 1 molecule C₃H₈

The Mass of a Mole

· 3 ways to measure mass:

1. Gram Atomic Mass (gam)

- The gam of an element is the mass (in grams) as shown on the periodic table for

an atom.

C = 12.01 g

2. Gram Molecular Mass (gmm)

o The mass of 1 mol of a molecular (covalent) compound.

SO3 = 1S + 3O = 32.10 g + 3(16.00 g) = 80.10 g

1. Gram Formula Mass (gfm)

o The mass of 1 mol of an ionic compound (also known as a formula unit).

CaI₂ = 1 Ca + 2 I = 40.10 g + 2(126.9 g) = 293.0 g

· We can not determine the gfm of CO₂ because it is a covalent compound. We need to determine the gmm.

Molar Mass

The mass in grams of 1 ml of substance.

1. Example:

How many grams ar in 9.45 moles of N₂O₃?

Known:

# moles N₂O₃ = 9.45 mole N₂O₃

1 mole N₂O₃ = 76 g N₂O₃

9.45 mol N₂O₃ x 76 g N₂O₃ = 718 g N₂O₃

1 mol 76 g N₂O₃

2. Example:

Calculate the mass of 2.50 mol Fe(OH)₂.

Known:

# moles Fe(OH)₂ = 2.50 mol Fe(OH)₂

1 mole Fe(OH)₂ = 89.965 g Fe(OH)₂

2.50 mol Fe(OH)₂ x 89.965 mol Fe(OH)₂ = 225 g Fe(OH)₂

1 mol Fe(OH)₂

3. Example:

Find the number of moles in 92.2 g Fe₂O₃.

Known:

Mass Fe₂O₃ = 92.2 g Fe₂O₃

1 mol Fe₂O₃ = 159.6 g Fe₂O₃

92.2 g Fe₂O₃ x 1 mol Fe₂O₃ = 0.578 mol Fe₂O₃

159.6 g Fe₂O₃

Volume of a Mole of Gas

· AT STP CONDITIONS (0°C AND 1 ATM OR 101 kPA) 1 MOLE OF GAS TAKES UP 22.4 L.

· For any gas at STP 1 mole contains 6.02 x 10²³ particles.

· Note that every gas at STP conditions will not have the same mass because each has its own molar mass based on its chemical formula.

1. Example:

What is the molar volume (in liters) of 0.60 mole SO₂ gas at STP?

Known:

# moles SO₂ = 0.60 mol SO₂

1 mol SO₂ = 22.4 L SO₂

0.60 mol SO₂ x 22.4 L SO₂ = 13 L SO₂

1 mol SO₂

· Density (g/L) at STP can be used to calculate molar mass.

1. Example:

The density of a gaseous compound containing carbon and oxygen is 1.964 g/L at STP. What is the molar mass?

Known:

Density = 1.964 g/L

1 mole at STP = 22.4 L

Molar mass = ? g/mol

1.964 g x 22.4 L = 44.0 g/mol

Percent Composition

· The relative amounts of each element in a compound is expressed as percent composition.

· I.e.: K₂CrO₄ => K = 40.3% ; Cr = 26.8% ; O = 32.9%

· Percentages should add up to 100%!

% mass of element = grams of element x 100%

grams of compound

1. Example:

An 8.20 g piece of Mg combines with 5.40 g of oxygen to form a compound. What is the percent composition.

Knowns:

Mass Mg = 8.20 g Mg

Mass O = 5.40 g O

Mass MgO = 13.60 g MgO

% Mg = 8.20 g Mg x 100% = 60.3 %

13.60 g MgO

% O = 5.40 g O x 100% = 39.7 %

13.60 g MgO

CHECK: 39.7 % + 60.3% = 100%

To Calculate Mass of 1 Mole of Compound

% mass = grams of element in 1 mol of compound x 100%

molar mass of compound

1. Example:

Calculate the percent composition of C₃H₈.

Known:

Mass C₃H₈ = 44.0 g/mol

Mass C in 1 mol C₃H₈ = 36.0 g

Mass H in 1 mole C₃H₈ = 8.0 g

% C = 36.0 g C x 100% = 81.8%

44.0 g C₃H₈

% H = 8.08 g H x 100% = 18.0 %

44.0 g C₃H₈

Percent as a Conversion Factor

1. Example:

What is the mass of carbon in 82.0 g of C₃H₈.

Known:

Mass C₃H₈ = 82.0 g C₃H₈

% by mass of C in C₃H₈ = 81.8% => in 100 g C₃H₈ there will be 81.8 g C

82.0 g C₃H₈ x 81.8 g C = 67.1 g

100 g C₃H₈

2. Example: Calculate the grams N in 125g of CO(NH₂)₂

Known:

Mass CO(NH₂)₂ = 125g

% N by mass =?

mass CO(NH₂)₂

a) % mass N = ____mass N____ x 100%

Mass CO(NH₂)₂

= 28.01g x 100% = 46.64% or 46.64 g in 100g CO(NH₂)₂

60.071g

b) 125g CO(NH₂)₂x ___46.64 N___ = 58.3g N.

100g CO(NH₂)₂

2.Example: Calculate the grams of N in 125 g NH₃.

Known: mass NH₃ = 125g

% N = ?

mass NH₃ = 17.04 g

% mass N = 14.01g x 100% = 82.22 or 82.22g N in 100g NH₃

17.04 g

125 g NH₃ x_82.22g N_ = 103 g N.

100 g NH₃

Calculating Empirical Formulas

*The empirical formula gives the lowest whole ratio of the atoms of the elements in a compound.

1. Example: What is the empirical formula of a compound that is 25.9% nitrogen and 74.1% oxygen?

Known:

% N = 25.9 % N

% O = 74.1 % O

Empirical formula = N_?O_?

a) In a 100g of compound there are 25.9 g of N and 74.1 f O.

25.9 N x 1 mol N = 1.85 mol N

14.0g

74.1g O x 1 mol O = 4.63 mol O

16.0g O

» N₁O_2.5 ß not correct because its not the lowest whole # ratio.

b) 1.85 mol = 1 mol N

1.85

4.63 mol O = 2.5 mol O

1.85

» N₁O_2.5 ß not yet

c) 1 mol N x 2 = 2 mol ß N₂O₅

2.5 mol O x 2 = 5 mol

2. Example: Calculate the empirical formula if we have 94.1 % O, 5.9% H.

a) á 94.1 O in 100g compound

á 5.9g H in 100g compound

ß 94.1g O x 1 mol N = 5.88 mol O

16.0g O

5.9g H x 1 mol H = 5.84

1.01g H

b) 5.88 mol O = 1.01 mol O

5.84

5.84 mol H = 1.0 mol H

5.84

» OH

Calculating Molecular Formulas

· Need to know the empirical formula and the molar mass in order to find the molecular formula.

1. Example: Calculate the molecular formula of the compound whose molar mass in 60.0g and empirical formula is CH₄N.

Known:

empirical formula = CH₄N

molar mass = 60.0g

molecular formula =?

a) Calculate empirical formula mass (efm).

empirical formula	efm	Molar mass/ efm	molecular mass
CH₄N	30.0	60.0/ 30.0 = 2	C₂H₈N₂

Please send comments and suggestions to Mr R

Name	Number	Weight

Melts		Boils