Wester Midterm Review

Info: This should cover everything on the study guide. Copy the notes from previous note packets to fill this out. Any unnecessary notes have been taken out.

* Replacement reactions are still in the notes even though we didn’t learn them as they are included in the study guide

Chapter 1: Matter and Change

MAKE STATES OF MATTER ITS OWN SECTION RIGHT BELOW THE MATTER TREE DIAGRAMTopics:

Definition of chemistry Classification of matter Properties of matter Energy

Definition of chemistry

Chemistry is a physical science, and is the study of…

chemical - any substance that has a definite composition (gold, glucose, water, CO)

Classification of matter

- the amount of three-dimensional space an object occupies.

- a measure of the amount of matter.

- anything that has mass and takes up space.

Basic Building Blocks of matter

pure substance – has a fixed composition; every sample has exactly the same characteristics and composition

element – a pure substance that cannot be broken down into simpler, stable substances and is made of one type of atom.

atom – the smallest unit of an element that maintains the chemical identity of that element.

compound – a substance that can be broken down into simple stable substances. Each compound is made from the atoms of two or more elements that are chemically bonded.

molecule – the smallest unit of a compound that maintains the chemical identity of that compound. (ionic ~ crystal/formula unit)

mixture – a blend of two or more kinds of matter, each of which retains its own identity and properties

homogeneous mixture – (solutions) uniform in composition (salt-water solution)

heterogeneous mixture – not uniform throughout (clay-water mixture)Separation techniques for mixtures:

___________________ - one substance caught, one goes through

___________________ - one substance vaporized, collected, and condensed

___________________ - one substance travels more slowly along a surface (paper)

Properties of matter

Physical Properties and Physical Changes

a characteristic that can be observed or measured without changing the identity of the substance, like melting point and boiling point.

a change in a substance that does not involve a change in the identity of the substance, like cutting, melting, and boiling.

mass, weight, length, volume, density, solubility (dissolving), state/phase, brittleness, flexibility, conductivity, energy, specific heat

States of matter

Matter

Description Solid Liquid Gas

particles touch each other

definite volume

definite shape

particles vibrate

particles have translational motion (place-to-place)particle arrangement

____________________ - a high-temperature physical state of matter in which atoms lose most of their electrons, particles that make up atoms.Chemical Properties and Chemical Changes

relates to a substance’s ability to undergo changes that turn it into different substances

chemical change = chemical reaction

carbon + oxygen carbon dioxide

possible evidence of a chemical change: color, temperature, odor, precipitate, bubbles

Demonstrations of physical and chemical changes:

Description Observations Physical or chemical?

Energy

Energy is involved when ___________________ or ___________________ changes occur.

Forms of energy:

Law of Conservation of Energy - energy can be or in a change, it is not destroyed or created

___________________ – average kinetic energy of the particles (motion)

___________________ – energy transferred between samples of matter because of a difference

in their temperatures

Chapter 2: Measurements and Calculations

Topics:1. Units of measurement2. Significant figures3. Scientific notation4. Useful chemistry measurements5. Analyzing lab data

Units of measurement

Le Système International d’Unités (SI) is a single set of measurements agreed upon by scientists around the world. It is similar to the metric system, but the “base unit” may be different.

SI units of measurement (Know the first 5)

Measurement SI unit Abbrev. Additional info

length meter m

mass kilogram kg Not grams

time second s

temperature Kelvin K Not C

amount of substance mole mol a way of counting particles

electric current ampere A Do not need to know.

luminous intensity candela cd Do not need to know.

Significant figures

Significant figures in a measurement are all the digits known with certainty plus the estimated (uncertain) digit

Tells you how many digits to round to in your answer. Tells you how precise your measurement tool is.

Atlantic-Pacific Method

Does the number have a decimal point?

Pacific Ocean Atlantic Ocean

P A

Present Absent

Start counting digits on the correct side of the number according to the A-P guide. Move to the first non-zero digit. Count that digit and every digit that comes after it, including any zeros.

Examples:

1. 40.0 mL = _____ significant figures

2. 982,500 cm = _____ significant figures

3. 20.540 mL = _____ significant figures

4. 0.06670 kg = _____ significant figures

5. 1.34 m = _____ significant figures

Calculations and significant figures

When adding or subtracting decimals, the answer must have the same number of digits to the right of the decimal point as there are in the measurement having the fewest digits to the right of the decimal point. (decimal places)

Example: 5.44 m - 2.6103 m =

For multiplication or division, the answer can have no more significant figures than are in the measurement with the fewest number of significant figures. (digits)

Example: 2.4 g/mL 15.82 mL =

Combining orders of operation

Follow order of operations, rounding the answer for each type of operation (add-sub or mult-div) along the way. (Depending on the textbook, you will get different instructions about this, but for this class, follow this method.)

Example: Find the density of a sample of an object that has a mass of 23.56 g. When placed in a graduated cylinder of 32.0 mL of water, the water level rises to 37.5 mL.

Scientific notation

Use the 2nd function EE combination on your graphing calculator instead of parentheses.

Examples:

1. 856,000 km = ___________________

2. 350. g = ___________________

3. 0.000433 g = ___________________

Useful chemistry measurements

mass – is a measure of the quantity of matter.

The SI standard unit for mass is the kilogram, however in class, grams (g), is used. Mass does not depend on gravity.

weight – (Derived, not SI) a measure of the gravitational pull on matter (Newtons, N).

volume – (Derived, not SI) the amount of space occupied by an object.

Multiply 3 lengths (length x width x height) The cubic centimeter, cm3, works well for solids. The milliliter (mL) or liter (L) works well for liquids. Displacement method: volume of object = final volume – initial volume

temperature – measure of the average kinetic energy of the particles in a sample

The SI standard unit for temperature is Kelvin, however lab thermometers are in C. K = C + 273

density – (Derived, not SI) the ratio of mass to volume, or mass divided by volume.

density = mass D = m Common units: g/mL or g/cm3

volume V

Density is a characteristic physical property of a substance that is useful for identification

Analyzing lab data

Accuracy & Precision

Accuracy refers to the closeness of measurements to the correct or accepted value of the quantity measured. (correct)

Precision refers to the closeness of a set of measurements of the same quantity made in the same way. (repeatable)

Interpreting data

average – add up the values and divide by the total number; round the answer to the correct number of significant figures

experimental value -

theoretical value -

percent yield = experimental value x 100 theoretical value

percent error = (experimental value – theoretical value) x 100 theoretical value

positive = too high negative = too low

Note: Depending on what you need to know about your data, you could make these calculations with data from each individual trial or you could average your data first and then make these calculations.

Example: A chemical reaction is run that theoretically releases 755 mL of gas at room temperature. A lab group repeats the experiment three times and collects the following volumes of gas: 780. mL, 730. mL, and 754 mL.

1. What is the average volume of gas produced?

2. Is their collection of data accurate according to the average?

3. Is their collection of data precise?

4. Calculate the percent yield based on the average.

5. Calculate the percent error based on the average.

6. Is the lab group’s average volume too high or too low compared with the theoretical amount of gas that should have been produced?

Chapter 3: Atoms: The Building Blocks of Matter

Topics:1. Foundations of atomic theory2. Development of the modern atomic theory3. The structure of the atom4. Counting subatomic particles5. The mole

Foundations of atomic theory

________________________________________________ - mass is neither created nor

destroyed during ordinary chemical reactions or physical changes

S + O2 SO2

H2 + O2 H2O

________________________________________________ - a chemical compound contains the

same elements in exactly the same proportions by mass regardless of the size of the sample or

source of the compound

________________________________________________ - if two or more different

compounds are composed of the same two elements, then the ratio of the masses of the second

element combined with a certain mass of the first element is always a ratio of small whole

numbers

Development of the Modern Atomic Theory

Dalton’s Atomic Theory Modern Atomic Theory

All matter is composed of extremely small

particles called atoms.

Atoms of a given element are identical in size,

mass, and other properties; atoms of different

elements differ in size, mass, and other propertie

Atoms cannot be subdivided, created, or

destroyed.

Atoms of different elements combine in simple

whole-number ratios to form chemical

compounds.

In chemical reactions, atoms are combined,

separated, or rearranged.

Discovery of the Electron and Nucleus

Scientist Time period

Contribution Sketch of experiment

J.J. Thomson late 1800s cathode ray tube

Robert Millikan late 1800s oil drop

Ernest Rutherford early 1900s gold foil

The structure of the atom

1. An atom is the smallest particle of an element that retains the properties of that element.

2. The nucleus is a small, dense region located at the center of an atom.

3. The nucleus is made up of at least one positively charged particle called a proton and

usually one or more neutral particles called neutrons.

4. Short-range nuclear forces hold the protons and neutrons together.

5. Surrounding the nucleus is a region with negatively charged particles called electrons.

6. A proton has a positive charge equal in magnitude to the negative charge of an electron.

7. Atoms are electrically neutral because they contain equal numbers of protons and

electrons.

8. A neutron is electrically neutral.

9. Atoms of different elements differ in their number of protons and therefore in the amount

of positive charge they possess.

10. The number of protons determines that atom’s identity.

11. Protons, neutrons, and electrons are often referred to as subatomic particles.

Counting subatomic particles

atomic number (Z) =

mass number =

isotopes =

The isotopes of a particular element all have the same number of protons and electrons

but different numbers of neutrons.

Most of the elements consist of mixtures of isotopes.

nuclide -

notations: uranium-235 or 92235 U

Example: How many protons and neutrons are in each nuclide?

612 C = 3

7 Li =

1224 Mg = 1

1 H =

Counting subatomic particles

protons (p+) =

neutrons (n0) =

mass number =

electrons (e-) =

Name Symbol Atomic number (Z)

Mass number

(A)

Protons Neutrons Electrons

iron 56

2040 Ca 40

28 31

argon-40 18

Atomic mass

relative atomic mass – the masses of elements are reported relative to the mass of carbon, which

has been arbitrarily assigned as the standard

atomic mass unit = amu = exactly 1/12 the mass of a carbon-12 atom

average atomic mass - weighted average of the atomic masses of the naturally occurring isotopes

of an element

The average atomic mass of an element depends on mass and relative abundance.

Example 1: Copper consists of 69.15% copper-63, which has an atomic mass of 62.929 601 amu, and 30.85% copper-65, which has an atomic mass of 64.927 794 amu.

The Mole

mole - amount of a substance that contains as many particles as there are atoms in exactly 12 g of

carbon-12.

Avogadro’s number = 6.02 1023 = number of particles in exactly one mole of a pure substance.

molar mass - mass of one mole of a pure substance (g/mol) (= atomic mass for an element)

Practice:

1. How many moles of silver, Ag, are in 3.01 1023 atoms of silver?

2. What is the mass in grams of 3.50 mol of the element copper, Cu?

3. What is the mass in grams of 1.20 108 atoms of copper, Cu?

Chapter 4: Arrangement of Electrons in Atoms

Topics:1. Properties of light2. Line-emission spectrum3. Bohr model4. Quantum model of the atom5. Atomic orbitals and quantum numbers6. Electron configurations and orbital diagrams7. Interpreting electron configurations

Properties of light

__________________________________________ – a form of energy that exhibits wavelike

behavior as it travels through space; moves at c = 3.00 x 108 m/s (light)

____________________ – () the distance between corresponding points on adjacent waves (m,

nm, cm); __________ (n) = 10-9

____________________ – () the number of waves that pass a given point in a specific time

(usually one second) (waves/second, hertz – Hz)

Equation connecting wavelength to frequency: c =

high frequency short wavelengthlow frequency long wavelength

__________________________________________ – electromagnetic radiation causes ejection of electrons from the surface of a metal

For a given metal, only specific frequencies eject electrons.

Max Planck proposed that energy is emitted in specific packets called quanta.

Albert Einstein proposed that light could be thought of as a stream of particles: photons.

Matter absorbs energy only in whole numbers of photons.

Planck’s constant = h = 6.626 x 10-34 JsE=hν=hc

λ

Line-emission spectrum

Since late 1800s, scientists knew that the light emitted by energized hydrogen was not continuous (all the colors), but limited to a few, specific frequencies.

Johan Balmer and Johannes Rydberg expressed the line spectrum for hydrogen mathematically.

line-emission spectrum – a series of specific wavelengths of light emitted from an energized sample of atoms (also called an atomic spectrum)

ground state – the lowest energy state of the electrons in an atom

excited state – the electrons of an atom have a higher potential energy than in the ground state

The energy diagram for hydrogen shows possible energy transitions for its electron as it emits energy from a higher energy level to a lower energy level.

Rydberg equation: ν=

RH

h ( 1nf

2−1ni

2 )Bohr model

Niels Bohr interpreted the data as electrons traveling in one of a limited number of specific orbits, or energy levels.

Since the electron in hydrogen would only emit specific frequencies of light, only specific amounts of energy were available to it.

The Bohr model represents the basic idea of principal energy levels very well and works for hydrogen

The Bohr model does not work for multi-electron atoms.

Quantum model of the atom

In classical mechanics, momentum = mass x velocity: p=mv

Louis de Broglie built on Einstein’s work and proposed that an electron behaves as a

particle and a wave.

deBroglie equations: momentum=p=h

λ wavelength=λ= h

mv

quantum theory – describes mathematically the wave properties of electrons and other

very small particles

Werner Heisenberg explained the difficulty in locating an electron that is both a particle

and a wave: the photons (light) we use to detect particles have about the same energy as

an electron so they knock an electron off its course.

Heisenberg uncertainty principle – it is impossible to determine simultaneously both the

position and velocity of an electron or any other particle.

Erwin Schrodinger developed equations to describe the most probable location of an

electron: Hyn=

1n2 ( me4

8 e02 h2 )Y n

wave function – mathematical equation that gives the probability of finding an electron

(90% of the time)

Matter AND radiation both have properties of waves AND particles

Atomic orbitals and quantum numbers

When solved, the Schrodinger equation provides quantum numbers that describe the electron and its orbital.

orbital – 3-D region around the nucleus that indicates the probable location of an electron.

quantum numbers – specify the properties of an atomic orbital and the electrons in the orbital

Information about quantum numbers

Name of quantum number

Possible values Description

Principal quantum number

Also: principal energy level shell

Provides a general sense of an electron’s overall energy

Higher values of n = further distance from the nucleus = more energy

Angular momentum quantum number

Indicates the shape of the orbital: s, p, d, f

Also: sublevel subshell

Magnetic quantum number

Indicates the orientation of an orbital around the nucleus.

Spin quantum number Accounts for the two types of magnetic fields created by electrons.

Using n + l describes the orbital.

Adding m and ms provides information about individual electrons.

Each orbital can hold 2 electrons.

Quantum numbers and number of electrons

Principal quantum number

(n)

Sublevels in principal energy

level(l = n-1)

Number of orbitals per

sublevel(-l thru +l)

Number of electrons per

sublevel

Number of electrons in the principal energy

level1

2

3

4

Electron configurations and orbital diagrams

An atom’s electron configuration shows the arrangement of its electrons, and there are three types of notation:

1. electron configuration notation – shows principal energy level, sublevel, and number of electrons in the sublevel (eg. 1s22s22p6)

2. orbital notation – shows all of the orbitals of a sublevel and the pairing of the electrons (uses arrows instead of superscripts)

3. noble gas configuration – abbreviates the electron configuration by using the prior noble gas ([Ar]3d54s1)

Rules of Electron Configurations and Orbital Diagrams

Rule 1:

Electrons fill orbitals of lower energy first.

The “s” sublevel is always the lowest energy sublevel within a principal energy level.

Rule 2:

Each orbital describes at most 2 electrons.

Electrons must have opposite spin (shown with up or down arrows).

Rule 3:

When electrons occupy a sublevel, one electron enters each orbital until all the orbitals

contain one electron with parallel spins, then they begin to pair up.

Aufbau diagram:

Practice electron configurations:

hydrogen, H (1 e-)

beryllium, Be (4 e-)

neon, Ne (10 e-)

Interpreting electron configurations

To identify an element by its electron configuration, add the superscripts, which equals the number of electrons. This number equals the atomic number for an atom.

Categorizing electrons:

______________________________ – electrons in the highest principal energy level; the outermost electrons; involved in bonding.

Using the electron configuration, add up the superscripts for the highest principal energy level only (1-8)

______________________________ – inner electrons held close to the nucleus; match the noble gas configuration; not usually involved in bonding

Practice: How many valence electrons do these elements have?

hydrogen helium

carbon chlorine

Connecting electron configuration to the periodic table

Practice: Identify these elements:

1. Electron configuration ends in 3s2. _____________________________________

2. Electron configuration ends in 5p5. _____________________________________

3. Electron configuration ends in 2p6. _____________________________________

Limits of the Aufbau diagram

The Aufbau diagram is a pattern that scientists came up with to model observations. The Aufbau diagram works best for atoms up to calcium (Z = 20) and for many after that. Transition metals have d orbitals that are very close to the valence s orbital, so for some

of them the Aufbau diagram does not correctly describe their arrangement. Electrons want the lowest energy possible with the most stability. They are affected by repulsions of nearby electrons and shielding. Examples:

chromium (Cr) [Ar]3d54s1

copper (Cu) [Ar]3d104s1

molybdenum (Mo) [Kr]4d55s1

silver (Ag) [Kr]4d105s1

gold (Au) [Xe]4f145d106s1

Chapter 5: The Periodic Law

Topics:1. History of the Periodic Table2. Sections of the Periodic Table3. Periodic trends

History of the Periodic Table

Dmitri Mendeleev

Noticed that when the elements were arranged in order of

increasing ___________________________, similarities

in their properties appeared at regular intervals.

Created a table in which elements with similar properties

were grouped together—a periodic table of the elements.

Predicted the existence and properties of

___________________ to fill empty spaces.

Henry Moseley

Discovered that the elements fit into patterns better when they

were arranged according to _____________________________,

rather than atomic mass.

_____________________________ - states that the physical and chemical properties of the elements are periodic, or systematic, functions of their atomic numbers.

The Modern Periodic Table

An arrangement of the elements by _____________________________ so that elements

with similar properties fall in the same column, or group.

Repeating patterns are referred to as _____________________________.

Sections of the Periodic Table

_______________ = column = up and down = family = similar chemical properties

_______________ = row = across = gradual changes in properties (trends)

Sections of the Periodic Table

Metals, nonmetals, and metalloids

s-block elements

Group 1 - _____________________________ – highly reactive

Group 2 - _____________________________ – a little less reactive, but do not exist

freely in nature

_____________________________ (1s1) does not share the same properties as the

elements of Group 1.

_____________________________ (2s2) is part of Group 18 because its highest

occupied energy level is filled by two electrons, which gives it helium chemical stability.

p-block elements

All of the elements of Groups 13–18 except _____________________________.

p-block and s-block elements are called the ___________________________________.

The properties of elements of the p block vary greatly.

________________________ – to the right, plus hydrogen

________________________ – boron, silicon, germanium, arsenic, antimony, tellurium

________________________ – to the left

Group 17 – halogens - most reactive nonmetals

Group 18 – noble gases – sublevels completely filled; extremely low reactivity

d-block elements

The d-block elements are metals with typical metallic properties and are often referred to

as .

f-block elements

Fit between Groups 3 and 4 in the sixth and seventh periods.

____________________ – top row - shiny metals similar in reactivity to the Group 2

alkaline metals

____________________ - bottom row - are all radioactive

Periodic Properties

All of these properties are most noticeable in the main group elements (blocks s and p).

Atomic radius

atomic radius -

Z = _____________________________ = _____________________________

Top to bottom: get larger more energy levels

Left to right: get smaller higher Z pulls electrons closer

Ionization Energy

______________ - an atom or group of bonded atoms that has a positive or negative charge.

______________ - any process that results in the formation of an ion

ionization energy, IE -

Top to bottom: decreases valence electrons further away from protons

Left to right: increases valence electrons closer to protons

Ionization energies are numbered:

IE1 = first ionization energy, IE2 = second ionization energy, etc.

Each successive IE is higher because the ion feels a stronger nuclear charge

_______________________________________ – the electrons between the nucleus and

the outer electrons reduce the attractive force felt by the outer electrons

Valence Electrons

valence electrons -

Electronegativity

electronegativity -

*Valence electrons hold atoms together in bonds and tend to concentrate near one of the atoms in the bond due to Z and atomic radius.

Top to bottom: decreases shared electrons are further from the pull of the protons

Left to right: increases more protons in the same principal energy level

Summary of periodic trends

Chapter 6: Chemical Bonding

Topics:1. Introduction to chemical bonding2. Covalent bonding and molecular compounds3. Molecular geometry4. Ionic bonding and ionic compounds

Section 1: Introduction to chemical bonding

Independent atoms have higher due to

electrostatic forces between protons and electrons.

Bonding potential energy for most atoms and creates more

Most atoms exist as part of a .

chemical bond -

ionic bond –

covalent bond –

Subtract electronegativities to predict bond type:

Bond type Electronegativity DIFFERENCE

Description

ionic

polar covalent

nonpolar covalent

Generalizations:

ionic = M + NM (metal + nonmetal)

covalent = 2 NM (two nonmetals)

Section 2: Covalent bonding and molecular compounds

molecule –

molecular compound –

chemical formula –

Remember: covalent = sharing = molecular

Formation of a Covalent Bond

The electrons of one atom and protons of a second atom one another.

The two nuclei and the electrons each other.

These two forces cancel out to form a at a length where the

potential energy is at a minimum.

When two atoms form a covalent bond, their shared electrons form

, this achieves a noble-gas configuration.

The Octet Rule

___________________ atoms generally do not react because their electron

configurations are stable.

Their outer s and p orbitals are completely _______________.

Other atoms can fill their outermost s and p orbitals by _______________ electrons.

octet rule –

Exceptions to the Octet Rule:

Some atoms cannot fit 8 electrons (_______________ is complete with 2 electrons;

_______________ has only 3 valence electrons and shares up to 6 electrons)

Some larger atoms can fit more than 8 electrons (_______________________________)

electron-dot notation –

Valence electrons

Electron-dot notation

Example Valence electrons

Electron-dot notation

Example

A pair of dots between two atoms represents a bond. Usually the pair of dots is turned into a straight line. A single bond is one line, a double bond is two lines, and a triple bond is three lines.

Lewis structures –

1. Write the electron-dot notation for each type of atom in the molecule. 2. Draw a skeleton structure. If C is present, it is the central atom. H and F are always

terminal atoms. The element with the lowest ionization energy/electronegativity goes in the middle (with some exceptions).

3. Count the total number of available valence electrons. If it is a negative ion, add the electrons; if it is a positive ion, subtract the electrons.

4. Count the total number of electrons needed for each atom to have a full valence shell.5. Find the number of bonding electrons: bonding = full – valence.6. Connect the atoms in the skeleton structure with pairs of bonding electrons. Usually a

straight line is used to represent the bond and it equals two electrons.7. Add the remaining electrons as lone pairs to each nonmetal atom (except hydrogen) to

complete the octet. lone pair electrons = valence – bonding 8. Count the electrons in the structure to be sure that the number of valence electrons used

equals the number available.

resonance structures –

Notes: Polyatomic ions: take into account the different numbers of electrons Hydrogen, H, has 2 valence electrons. Boron, B, has 6 valence electrons.

Some covalent bonds do NOT form molecules

network covalent solids –

Can be elements or compounds: carbon silicon, silicon dioxide

Section 3: Molecular Geometry

The properties of molecules depend on the bonding and the

Molecular geometry and the polarity of individual bonds determines the overall

_______________ of the molecule, which in turn helps predict its properties.

hybridization –

Example: Carbon is 1s22s22p2. Based on configuration, it seems as if carbon should have one

unshared pair of electrons and two lone electrons for bonding, but instead it has 4 bonding sites.

VSEPR theory – (“VES-pur” = “valence-shell electron-pair repulsion”)

The name of the shape of a molecule refers to the positions of _______________ only.

Molecular geometries

A = central atom

B = bonded atom of any kind

E = electron pairs

Shape ABE notation Sketch

AB2

AB2E or AB2E2

AB3

AB3E

AB4

Intermolecular Forces

intermolecular forces –

The boiling point of a liquid is a good measure of the intermolecular forces between its

molecules: the _______________ the boiling point, the _______________ the forces

between the molecules.

Intermolecular forces vary in strength but are generally _______________ than bonds

between atoms in molecules, ions in ionic compounds, or metal atoms in solid metals.

dipole –

dipole-dipole forces –

hydrogen bonding –

(hydrogen bonds have FON)

London dispersion forces –

Contributing to strength:

Hydrogen bonding

Increases with more electrons, larger molecular weight, and greater surface area

Section 4: Ionic bonding and ionic compounds ionic compound –

formula unit –

Most ionic compounds exist as ______________________________.

A crystal of any ionic compound is a

of positive and negative ions mutually attracted to each other.

The chemical formula of an ionic compound represents not molecules, but the

(formula unit)

Formation of Ionic Compounds

The _______________ electronegativity of nonmetals means they will gain an electron

and take on a _______________ charge.

The _______________ electronegativity of metals means they will lose an electron and

take on a _______________ charge.

crystal lattice –

lattice energy –

polyatomic ion –

Chapter 7: Formulas and Compounds

Topics Section 1 - Ionic names and formulas Section 2 - Molecular names and formulas Section 4 - Mole & mass conversions Section 5 - Percent composition Section 6 - Determining chemical formulas

Section 1 - Ionic names and formulas

IONIC COMPOUNDS – the chemical formula represents a “formula unit,” which is the lowest whole-number ratio of ions in the lattice.

Monatomic ions – formed when atoms gain or lose electrons

Groups: 1+ 2+ 3+ 3- 2- 1-

Metals with unpredictable charges (form cations)

Some d-block elements and metals below the staircase can form more than one ion. You

will be able to figure out the charge from the name or formula that it is part of. However,

it is important to recognize when these are being used so that they are named correctly:

o See the elements on your reference sheet that have a dotted line in their boxes.

It is difficult to predict the charges that some of these metals take on. MEMORIZE

THESE TWO: zinc: Zn2+ AND silver: Ag+

Polyatomic ions – Re-memorize the list from the reference sheet.

Ionic formulas:1. Write the symbol with the charge.2. Criss-cross the charges to find subscripts.3. Drop the + or – symbol.4. Do not write the number 1.5. Reduce the subscripts, if necessary.6. Use parentheses around polyatomic ions with a subscript larger than 1.

Naming:7. Write the regular name of the positive (metallic) ion.8. Use the Roman numeral designation after the name of metals with multiple charges.9. Change the ending of the negative (nonmetallic) ion to –ide.10. Use the name of polyatomic ions as is.

Hydrates – water makes up part of the crystal structure

barium hydroxide octahydrate Ba(OH)28H2Olead (II) perchlorate trihydrate Pb(ClO4)23H2O(You will not have to name these. They are here so that you are familiar with them.)

Section 2 - Molecular names and formulas

MOLECULAR FORMULAS – provide the exact number and kind of atoms in the molecule. There can be multiple combinations between the same two elements, so you have to be given either the name or the formula. Do not reduce the subscripts.

1. If given the name, use the prefixes to write the subscripts.2. If given the formula, use the subscripts to write the prefixes.3. Change the ending of the second element to –ide.4. Do not write the prefix mono for the first element.5. The o or a at the end of a prefix is usually dropped if the name starts with a vowel.

Prefixes1 mono 6 hexa2 di 7 hepta3 tri 8 octa4 tetra 9 nona5 penta 10 deca

Special cases. Memorize:

water = H2O (not dihydrogen monoxide)

ammonia = NH3 (not nitrogen trihydride)

Diatomic elements - Seven elements exist as bonded as pairs. Memorize:

hydrogen – H2

nitrogen – N2

oxygen – O2

fluorine – F2

chlorine – Cl2

bromine – Br2 (liquid)

iodine – I2 (solid)

Acids and salts. Memorize:

Binary Acids Oxyacids

(hydro – “root” – ic acid)

HF – hydrofluoric acid HNO3 – nitric acid

HCl – hydrochloric acid HNO2 – nitrous acid

HBr – hydrobromic acid H2SO4 – sulfuric acid

HI – hydroiodic acid H2SO3 – sulfurous acid

H2S – hydrosulfuric acid H3PO4 – phosphoric acid

HCN – hydrocyanic acid CH3COOH – acetic acid

Section 4 - Mole and mass conversions

Chemical formulas are a connection between the lab and chemistry at the particle level.

mole = Avogadro’s number = _______________________________________

molar mass –

The formula mass adds up to the same number, but is in units of amu. For this section, practice calculating molar mass, mole – mass, and mole – particle

conversions.

Section 5 - Percent composition

percent composition –

% of element in the compound =

Calculate directly from lab measurements by dividing for percentage. Calculate from a chemical formula starting with mole conversions. This is based on the

assumption that you have one mole of the substance.Section 6 - Determining chemical formulas

empirical formula –

Two molecular formulas: ethane: C2H4 or butene: C4H8

Both have an empirical formula of CH2. Starting measurements are either lab measurements of mass or percent composition.

Example: The percent composition of ammonia is 82.4% nitrogen and 17.6% hydrogen. Calculate the empirical formula of ammonia.

Step 1: Assume the sample is 100 g and make each percentage become a mass.

Step 2: Convert mass to moles.

Step 3: Divide each mole value by the smallest mole value available.

Step 4: Round each division problem to the nearest whole number and use them as subscripts

molecular formula –

Example: Fructose has an empirical formula of CH2O and a molar mass of 180.00 g/mol. What is its molecular formula?

Step 1: Calculate the formula mass of the empirical formula.

Step 2: Find the multiplying factor, x.

x = molecular formula massempirical formula mass

Step 3: Multiply each subscript by the whole number multiplying factor.

Chapter 8: Chemical Reactions

Topics1. Describing chemical reactions2. Skeleton equations3. Balanced chemical equations4. Synthesis reactions5. Decomposition reactions6. Single replacement reactions7. Double replacement reactions8. Combustion reactions

Describing chemical reactions

Possible indications of a chemical reaction:

1. ______________________________________________ released

2. ____________ produced

3. ________________________ formed

4. ____________ change

chemical equation -

Characteristics of chemical equations:

1. All and must be identified.

2. The reactants and products must be written with the correct

3. The must be

satisfied (balancing).

Levels of writing a chemical equation:

WORD: methane + oxygen carbon dioxide + water

SKELETON: CH4(g) + O2(g) CO2(g) + H2O(g) (not balanced)

BALANCED: CH4(g) + 2O2(g) CO2(g) + 2H2O(g)

Skeleton equations

skeleton equation -

Memorize this list:

Diatomic elements:

S8, P4 (solids)

_________________________________________________ are generally solids.

Liquid elements:

Gases: and all

Ionic compounds are usually .

Acids are usually .

Symbol Explanation

separates two reactants or two products

“yields” – separates reactants from products

reversible reaction

solid

liquid

gas

aqueous, dissolved in water

gas product

precipitate formed

Heat is supplied to the reaction

The formula of a catalyst is written above the yield symbol (in this example, platinum, Pt)

Balanced chemical equations

balanced equation –

Balancing equations demonstrates the

(mass is neither created nor destroyed in a closed system)

Use coefficients to balance the equation (not subscripts).

6 H2O“6” (coefficient) refers to total number of particles “2” and “1” (subscripts) refer to the number

of atoms in the substance and are part of its identity

Guidelines for balancing equations:1. This is done by inspection and trial-by-error, but mini charts can help.2. Balance the different types of atoms one at a time.3. First balance the atoms of elements that are combined and that appear only once on each

side of the equation.4. Balance polyatomic ions that appear on both sides of the equation as single units.5. Balance H and O atoms last.6. Count to make sure each type of atom is balanced.

Synthesis Reactions

synthesis reactions -

A + B AB

Metal + nonmetal: one ionic product

Na(s) + O2(g)

Nonmetal + oxygen: one molecular product

C(s) + O2(g)

H2(g) + O2(g)

Metal oxides + water: one ionic product ending with hydroxide (OH)

CaO(s) + H2O(l)

Decomposition Reactions

decomposition reactions -

CD C + D

Binary compounds: two elements

H2O(l)

Metal carbonates: metal oxide + carbon dioxide

CaCO3(s)

Metal hydroxides: metal oxide + water

Ca(OH)2(s)

Metal chlorates: metal chloride + oxygen

KClO3(s)

Single Replacement/Displacement

single replacement -

E + FG F + EG or E + FG G + FE

**Use the Activity Series to predict whether or not a reaction will happen.

Metal plus ionic compound

Al(s) + Pb(NO3)2(aq)

Metal plus acid: ionic compound + hydrogen

Mg(s) + 2HCl(aq)

Halogen plus ionic compound

Cl2(g) + KBr(aq)

Double Replacement/Displacement

double replacement -

HI + JK HK + JI

**The reaction will happen if a precipitate, gas, or water is formed.

Formation of a precipitate

PbI2(s) + KNO3

Formation of a gas

FeS(s) + HCl(aq)

Formation of water

HCl(aq) + NaOH(aq)

Combustion

combustion -

CxHy + O2(g) CO2(g) + H2O(g/l)

combustion of methane

CH4(g) + O2(g)