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9/6/12
What is chemistry?
A chemical is any substance that has a definite composition.
A chemical reaction is the process by which one or more substances change to
produce one or more new substances.
Physical states of matter
-The states of matter are the physical forms of matter which are solid, liquid, gas,
and plasma.
Macroscopic refers to what you can see with the unaided eye
Microscopic refers to what you would see if you could see individual atoms.
9/7/12
Solids have a fixed volume and shape that result from the way their particles are
arranged.
Liquids have a fixed volume but not a fixed shape
Gases have neither a fixed volume nor shape
Physical Changes are changes in which the identity of a substance doesn’t change
- Changes of state are physical changes
Chemical Changes occur when the identities of substances change and new
substances form.
9/10/12
Chemical Changes
Mercury(II) oxide  mercury + oxygen
Reactants are the substances on the left-hand side of the arrow
(they are used up in the reaction
Products are the substances on the right-hand side of the arrow
(they are produced in the reaction)
Evidence of a chemical change
- the evolution of a gas
- the formation of a precipitate (solid substance forms and typically sinks to the
bottom)
- the release or absorption of energy
- a change in temperature or the giving off of light
- a color change in the reaction system
Matter
Matter- Peanut butter, water, fish, garbage, human brain, carbon dioxide, air,
yourself, tree
Not Matter- light, time, motion, an idea, energy
Distinguish between different characteristics of matter, including mass, volume, and
weight.
Identity and use SI units in measurements and calculations
Set up conversion factors, and use them in calculations
Identify and describe physical properties, including density
Identify chemical properties
Matter had Mass and Volume
- Matter is anything that has mass and volume
- Volume is the space an object occupies
- Mass is the quantity of matter in an object
o Devices used for measuring mass in a laboratory are called balances or
scales
- Weight is the force produced by gravity acting on a mass
9/11/12
Units of Measurements
- when working with numbers, be careful to distinguish a quantity and its units
o Quantity describes something that has magnitude, size, or amount
o Unit is a quantity adopted as a standard of measurement
Meter
kilogram
second
ampere
Unit
symbol
m
kg
s
A
kelvin
K
candela
cd
mole
mol
Unit name
Name
Symbol
Factor
Prefix
exa
peta
tera
giga
mega
kilo
hecto
deca
Symbol
E
P
T
G
M
K
h
da
deci
centi
milli
micro
nano
pico
d
c
m
μ
n
p
100
1000m
10006
10005
10004
10003
10002
10001
10002/3
10001/3
10000
1000-1/3
1000-2/3
1000-1
1000-2
1000-3
1000-4
decada
101
Quantity name
length
mass
time
electric current
thermodynamic
temperature
Dimension
symbol
L
M
T
I
Quantity symbol
l (a lowercase L), x, r
m
t
I (an uppercase i)
T
Θ
luminous intensity
Iv (an uppercase i with lowercase
non-italicized v subscript)
J
amount of substance
n
N
hectoh
102
kilo- megak
M
103
106
giga- teraG
T
109
1012
10n
Decimal
1018
1000000000000000000
1015
1000000000000000
1012
1000000000000
109
1000000000
106
1000000
103
1000
102
100
101
10
100
1
10-1 0. .1
10-2 0. .01
10-3 0. .001
10-6 0. .000001
10-9 0. .000000001
10-12 0 .000000000001
peta- exa- zettaP
E
Z
1015
1018 1021
Short scale
quintillion
quadrillion
trillion
billion
billionth
trillionth
Long scale
trillion
billiard
billion
milliard
million
thousand
hundred
ten
one
tenth
hundredth
thousandth
millionth
milliardth
billionth
A conversion factor is a simple ration that relates two units that express a measurement of
the same quantity
you can construct conversion factors between kg and grams as follows:
1kg/1000g or 1000g/1kg
Given: 4.5kg
?g
4.5kg 1000g = 4500g
1kg
-
Given: 208lb
? kg
208lb .45kg = 93.6kg
1lb
9/13/12
Given: 0.851L
?mL
.851L 1000mL = 851.0mL
1L
Derived Units
- Many quantities you can measure need units other than the seven basic SI
units
- These units are derived by multiplying or dividing the base unit
o Speed is distance divided by time. A derived unit of speed is
meters/second (m/s)
o A rectangle’s area is found by multiplying its length (in meters) by its
width (also in meters)
 Its unit is square meters (m2)
- Volume is another commonly used derived unit
o The volume of a book can be found by multiplying its length, width,
and height.
o The unit of volume is the cubic meters (m3)
o This unit is too large and inconvenient in most labs. Chemists usually
use the liter (L)
 1L = 1000mL = 1000cm3
 1mL = 1cm3
Properties of Matter
Physical Properties
- A physical property of a substance is a characteristic that does not involve a
chemical change.
- Physical properties of a substance can be determined without changing the
nature of a substance.
- Physical properties include malleability, color, texture, state, density, melting
point, and boiling point.
9/14/12
Properties of Matter
Density is the Ratio of Mass to Volume
-
The density of an object is the mass if the object divided by volume of the
object
Densities are expressed in derived unties such as g/cm3 or g/mL
Density is calculated as follows:
o Density = mass/volume OR D = m/v
The density of a substance is the same no matter what the size of the sample
is.
Because the density of a substance is the same for all samples, you can use
this property to help identify substances.
9/18/12
Do Now
Kilo = 1000
.01 = centi
45g = .045 kg
100mL = .1L
4.8cg = 48mg
Properties of Matter
Chemical Properties
- A chemical property is a property of matter that describes a substance’s ability
to participate in chemical reactions – i.e.- Reactivity
- A chemical property of many substances is their reactivity with oxygen.
o Rusting, corrosion
- Some substances break down into new substances when heated
Classifying Matter
- An atom is the smallest unit of an element that maintains the properties of that
element.
- Matter exists in many different forms but there are only 118+ types of atoms.
- Atoms are joined together to make up all the different kinds of matter.
Pure Substance
- A pure substance is a sample of matter, either a single element or a single
compound, which has definite chemical and physical properties.
- Elements are pure substances that only contain one kind of matter (atom).
They cannot be separated or broken down into simpler substances by ordinary
chemical means. (exception- Higgs Boson particle)
o Each element has its own unique set of physical and chemical
properties
- A molecule is the smallest unit of a substance that keeps all of the physical
and chemical properties of that substance.
- A molecule usually consists of two or more atoms combined in a definite ratio
o Except diatomic elements
- Diatomic elements exist as two atoms of the same element joined together.
o H2 O2 N2 F2 Cl2 Br2 I2
o Acronym- HOFBrINCl
- Some elements, such as oxygen, phosporus, sulfur, and carbon, have many
molecular forms
- An allotrope is one of a number of molecular forms of an element
-
The properties of allotropes vary widely
o Carbon- graphite, diamond, buckyballs
9/20/12
Pure Substances
Some Elements Exist in More than One Form
- Oxygen exists as allotropes
- Oxygen gas (O2) is colorless and odorless
- Ozone (O3) is toxic and pale blue
Compounds are Pure Substances
- Pure substances that are not elements are compounds. Compounds are
composed of more than one kind of atom.
o Example: carbon dioxide
- There may be easier ways of preparing them, but compounds can be made
from their elements.
- Compounds can be broken down into their elements, often with difficulty.
Compounds are Represented by Formulas
- Compounds have characteristic in properties and composition
- Compounds can be represented by an abbreviation or Formula
o A formula has subscripts, which represent the number of different
atoms in the compound.
o Example: H2O has 2 hydrogen atoms and 1 oxygen atoms
- Molecular formulas give information only about what makes up a compound.
o Example: the molecular formula for aspirin is C9H8O4
- A structural formula shows how the atoms are connected
o Two-dimensional models do not show the molecule’s true shape
9/21/12
How is matter classified?
- A call and stick model shows the distances between atoms and the angles
between them in three dimensions.
- A space filling model attempts to represent the actual size of the atoms nd not
just their relative position
- These models convey different information
Mixtures
- A mixture is a combination of two or more substances that are not chemically
combined
- Air is a mixture of mostly nitrogen and oxygen
- Water is not a mixture
o The H and O atoms are chemically bonded
o The ratio of H to O atoms is always 2 to 1
- The properties of the mixture may vary
- An alloy is a solid mixture
o Example: an alloy of gold and other metal atoms is stronger than pure
gold
 18-karot gold contains 18 grams of gold per 24 grams of alloy
 14-karot gold contains 14 grams of gold per 24 grams of alloy
Homogenous mixture
- A homogenous mixture describes something that has a uniform structure or
composition throughout.
o Example: gasoline, syrup, and air
- Homogenous mixtures have the same properties throughout
Heterogeneous Mixtures
- A heterogeneous mixture describes something that is composed of dissimilar
components
o Example: a mixture of sand and water
- Any two samples of a heterogeneous mixture will have the different
proportions of ingredients
o Heterogeneous mixtures have different properties throughout
- Mixtures can be either homogenous or heterogeneous
9/24/12
Do Now
37 gigabytes = 37,000,000 kilobytes
763 mL = .763L
45 kilograms = 45,000,000 mg
Understanding Concepts
Tea is best classified as a homogenous mixture
He is an element
CHAPTER 2
Objectives
- Explain that physical and chemical changes in matter involve transfers of
energy
- Apply the law of conservation of energy to analyze changes in matter
- Distinguish between heat and temperature
- Convert between Celsius and Kelvin
Energy and Change
- Energy is the capacity to do work, such as moving an object, forming a new
compound, or generating light (Electromagnetic Energy)
- Energy is always involved when there is a change of matter
Changes in Matter Can Be Physical or Chemical
- Ice melting and water boiling are examples of physical changes
- A Physical Change is a change of matter that affects only the physical
properties of matter.
- In contrast, the reaction of hydrogen and oxygen to produce water is an
example of a chemical change
- A Chemical Change is a change that occurs when one or more substances
change into entirely new substances with different properties
- A chemical change occurs whenever a new substance is made
- All physical and chemical changes involve a change in energy
- Sometimes energy much be supplied for the change in matter to occur
- Evaporation is the change of a substance form a liquid to a gas due to the
absorption of Energy
9/25/12
Energy and Change
Endothermic and Exothermic Processes
- Any change in matter in which energy is absorbed from the surrounding in an
Endothermic process
o The melting of ice and boiling of water are examples of physical
changes that are Endothermic
o As the chemicals react, energy is absorbed. Energy is a Reactant
- Any change in matter in which energy is released is an Exothermic process
o The freezing of water and condensation of water vapor are two
examples of physical changes that are exothermic processes.
o Energy is released. Energy is a Product.
- Energy can be absorbed (Endothermic) by the surroundings or released
(exothermic) to the surroundings, but it cannot be created or destroyed.
- The Law of Conservation of Energy states that during any physical or
chemical change, the total quantity of energy remains constant.
- In other words, energy cannot be created nor destroyed.
Energy is often transferred
- A system consists of all the components that are being studied at any given
time.
o The chemicals are the system, not the beaker they are in.
- The surroundings include everything around outside the system.
- Energy exists in different forms, including:
o Heat, Light, Potential-Chemical, kinetic-Motion, Sound, Electrical,
Pressure waves
- The transfer of energy between a system and its surroundings can involve any
one of these forms of energy
Heat
- Heat is the energy transferred between objects that are different temperatures.
- Heat energy is always transferred from a warmer object to a cooler object
- For example, when ice cubes are placed in water, heat energy is transferred
from the water to the ice.
- Energy is also transferred as heat during chemical changes.
9/27/12
Energy Can Be Absorbed as Heat
- In an endothermic reaction, energy is absorbed by the chemicals that are
reacting.
Heat is Different from Temperature
- Scientists define Temperature as a measurement of the average kinetic energy
of the random motion of the particles in a substance
- The transfer of energy as Heat can be measured by calculated changes in
Temperature
Temperature is Expressed Using Different Scales
- Thermometers are usually marked with the Fahrenheit or Celsius temperature
scales
- A third temperature scale uses the unit Kelvin, K
The zero point of the Celsius scale is designed as the freezing point of water
The zero point on the Kelvin scale is designated as absolute zero, the
temperature at which the minimum average kinetic energies of all particles
occur.
- There are no negative temperatures in Kelvin
Transfer of Heat May Not Affect Temperature
- The transfer of energy as heat does not always result in a change of
temperature
o For example, consider that happens when energy is transferred to a
mixture of ice and water
o As energy is transferred as heat to the ice-water mixture, the ice cubes
start to melt.
o The temperature of the mixture remains at 0C until all the ice has
melted.
-
Transfer of Heat Affects Substances Differently
- The specific heat of a substance is the quantity of energy as heat that mist be
transferred to raise the temperature of 1g of a substance 1K
- The SI (metric) Unit for energy is Joule (L)
- Specific Heat is expressed in joules per gram Kelvin (J/gK)
- Metals tend to have low specific heats
- Water has a high specific heat
9/28/12
Scientific Method
- The Scientific Method is a series of steps followed to solve problems,
including:
o Collecting data
o Formulating a hypothesis
o Testing the hypothesis
o Stating conclusions
- An experiment is the process by which scientific ideas are tested.
- A Hypothesis is a reasonable and testable explanation for observations.
- A Variable is a factor that could affect the results of an experiment.
- When a variable is kept constant form one experiment to the next, the variable
is called the control.
- The procedure is called a Controlled Experiment
- Any hypothesis that withstands repeated testing may become part of a theory
- In science, a Theory is a well-tested explanation of observations
Theories and Laws Have Different Purposes
- Some facts in science always hold true. These facts are called laws.
- A Law is a statement or mathematical expression that reliably describes a
behavior of the natural world.
- A theory is an attempt to explain the cause of certain events in the natural
world.
- For example, the Law of Conservation of Mass states that the products of a
chemical reaction have the same mass as the reactants have.
- Keep in mind that a hypothesis predicts an event, a theory explains it, and a
law describes it.
Models Can Illustrate the Microscopic World of Chemistry
- A Model represents an object, a system, a process, or an idea
10/3/12
Measurements and Calculations in Chemistry
Accuracy and Precision
- The Accuracy of a measurement is how close the measurement is to the true
or actual value
- Precision is the exactness of a measurement
- It refers to how closely several measurements of the same quantity made in
the same way can agree.
Significant Figures
- Scientists report values using significant figures
- The Significant Figures of a measurement or a calculation consists of all the
digits known with certainty as well as one estimated, or uncertain, digit.
- The last digit or significant figure reported after a measurement is uncertain
and estimated
Rules for Determining Significant Figures
- Nonzero digits are always significant
- Zeros between nonzero digits are significant
- Zeros in front of nonzero digits are not significant
- Zeros both at the end of a number and to the right of a decimal point are
significant
- Zeros both at the end of a number but to the left of a decimal point may not be
significant. If a zero has not been measured or established, it is not significant.
A decimal point placed after zeros indicated that the zeros are significant
10/8/21
Do Now
Define: Density- mass divided my volume
Formula- D=m/v
Units- g/mL or g/cm3
Significant Figures
- To determine the number of significant figures in the answer, you must first
find the number of significant figures in the values used to calculate the
answer.
When multiplying and dividing, the answer cannot have more sigfigs than the
measurement with the smallest number of sigfigs.
- With addition and subtraction, the result can be no more certain than the least
certain number in the calculation.
- Conversion factors and countable values have unlimited numbers of sigfigs
OR it does not limit my calculation.
10/9/12
Significant Figures
Example: A student heats 23.62g of a solid and observes that the temperature
increases from 21.6C to 36.79C. Calculate the temperature increase per
gram of solid.
36.79C - 21.6C = 15.19C = 15.2C
15.2C/23.62g = 0.644C/g
Specific Heat Depends on Various Factors
- Specific heat depends on the nature of the material that is changing
temperature, the mass of the material, and the size of the temperature change.
- Recall that the specific heat is the quantity of energy that must be transferred
as heat to raise the temperature of 1g of a substance by 1K.
- The specific heat (cp) of a substance at a given pressure (p) is calculated by
the following formula
o Specific heat = cp = q / mxT (q-heat in joules, m-mass in grams,
T- change in temperature in K)
Example: A 4.0g sample of glass was heated from 274K to 314K and was found
to absorb 32J of energy as heat. Calculate the specific heat of this glass.
cp = q
= 32J
= .20J/gK
mxT (4.0g)(40K)
10/10/12
Scientific Notation
- Very large and very small numbers are often written in scientific notation.
- A number in scientific notation has 2 parts
- The first part is a number that is between 1 and 10
- The second part consists of a power of 10
- To write the first part of the number, move the decimal to the right of the left
so that only one nonzero digit is to the left of the decimal.
Scientific Notation with SigFigs
- Use scientific notation to eliminate all place-holding zeros
- Move the decimal in an answer so that only one digit is to the left, and change
the exponent accordingly. The final value must contain the correct number of
sigfigs.
o 1001000000 = 1.001x109
o 0.0000456 = 4.56x10-5
o 0.0000036mm = 3.6x10-6mm
o 1450000mg = 1.45x106mg
-
10/11/12
SigFig Practice
1890.01m – 6 sigfigs
0.01702L – 4 sigfigs
Specific heat- is the quantity of energy that must be transferred as heat to raise the
temperature of 1g of a substance by 1K.
Specific heat = cp = q / mxT (q-heat in joules, m-mass in grams, T- change in
temperature in K)
10/15/12
CHAPTER 3
3.1 Atomic Theory
- The idea of an atoms theory is more than 2000 years old.
- Until recently, scientists had never seen evidence of atoms.
- The law of definite proportions, the law of conservation of mass and the law
of multiple proportions support the current atomic theory
The Law of Definite Proportions
- The Law of Definite Proportions states that a chemical compound always
contains the same elements in exactly the same proportions by weight or
mass.
- The law of definite proportions also states that every molecule of a substance
is made of the same number and types of elements.
The Law of Conservation of Mass
- The law of conservation of mass states that mass cannot be created nor
destroyed in ordinary chemical and physical changes.
- The mass of the reactants is equal to the mass of the products.
The Law of Multiple Proportions
- The law of multiple proportions states that when two elements combine to
form two or more compounds, the mass of one element that combines with a
given mass of the other is in the ratio of small whole numbers.
Dalton’s Atomic Theory
- In 1808, John Dalton developed an atomic theory.
- Dalton believed that a few kinds of atoms made up all matter.
- According to Dalton, elements are composed of only one kind of atom and
compounds are made from two or more kinds of atoms.
- Dalton’s Theory Contains Five Principles
o All matter is composed of extremely small particles called atoms,
which cannot be subdivided, created, or destroyed.
o Atoms of a given element are identical in their physical and chemical
properties.
o Atoms of different elements differ in their physical and chemical
properties.
o Atoms of different elements combine in simple, whole number ratios
to form compounds.
o In chemical reactions, atoms are combined, separated, or rearranged
but never created, destroyed, or changed.
Data gathered since Dalton’s time shows that the first two principles are not
true in all cases.
HW- Read Section 3.1, do 1-9 (pg 78)
10/16/12
3.2 Structure of Atoms
Subatomic Particles
- Experiments by several scientists in the mid-1800’s led to the first change to
Dalton’s atomic theory.
- The smaller parts that make up atoms are called subatomic particles
- The three subatomic particles that are most important for chemistry are the
electron, proton, and neutron.
Electrons were discovered using Cathode Rays
- JJ Thompson studied currents
- Thomson observed a glowing beam that came out of the cathode and struck
the anode and the nearby glass walls of the tube.
o He called these rays cathode rays
o The glass tube Thomson used is known as a cathode-ray tube.
An electron has a Negative Charge
- Thomson’s experiments showed that a cathode ray consists of particles that
have mass and a negative charge
- There particles are called electrons
- An electron is a subatomic particle that has a negative electric charge
- Electrons are negatively charged, but atoms have no charge
o Atoms contain some positive changes that are balanced by negative
charges.
- Electron: e, e-, or –10e ; charge of 1.602x10-19C ; common charge notation –1 ;
mass is 9.109x10-31kg
Rutherford Discovered the Nucleus
- Thomson proposed that the electrons of an atom were embedded in a
positively charged ball of matter. His model of an atom was named the plumpudding model.
- Ernest Rutherford performed the gold foil
experiment, which discovered the nucleus of an
atom.
o A beam of small, positively charged
particles, called alpha particles, was
directed at a thing gold foil.
- Rutherford called the space around the atom, the
nucleus.
10/17/12
- The nucleus is the dense, central portion of the atom.
- The nucleus is made up of protons and neutrons
- The nucleus has all the positive charge, and nearly all of the mass, but only a
very small fraction of the volume of the atom.
-
Protons
- Protons are the subatomic particles that have a positive charge and that is
found in the nucleus of an atom.
o The number of protons of the nucleus is the atomic number, which
determines the identity of an element.
o Because protons and electrons have equal but opposite charges, a
neutral atom must contain equal numbers of protons and electrons.
- Neutrons are the subatomic particles that have no charge and that are found in
the nucleus of an atom
Charge
Particle
Relative Charge
Mass
Relative mass
Proton
+1.60 x 10-19 C
+1
1.672 x 10-24 g
1 amu
Neutron
neutral
0
1.674 x 10-24 g
1 amu
Atomic Number
- The number of protons that an atom has is known as the atomic number.
o The atomic number is the same for all atoms of an element.
o No two elements can have the same atomic number.
- Atomic numbers are always whole numbers.
- The atomic number also reveals the number of electrons in an element.
o For atoms to be neutral, the number of negatively charged electrons
must equal the number of positively charged protons.
10/18/12
Do Now
1) Who discovered the electron? JJ Thomson
How? Cathode Ray Tube
When? Early 1900’s
2) What did Rutherford determine? Atom is made up of mostly space and there
is dense nucleus that is positively charged in the 1900’s
3) Who developed the first Atomic Theory? Dalton in the early 1800’s
Mass Number is the Number of Particles in the Nucleus
- The mass number is the sum of the number of protons and neutrons of an
atom.
o Mass# = p+ + n0
o Mass of e- = 0 (charge –1)
o Mass of p+ = 1 (charge +1)
o Mass of n0 = 1 (charge 0)
- You can calculate the number of neutrons in an atom by subtracting the
atomic number (the number of protons) from the mass number (the number of
protons and neutrons)
o Atomic# - p+
o Mass# - p+ + n0
o # of neutrons = mass# - atomic#
- Unlike the atomic number, the mass number can vary depending on the
number of neutrons.
Example: a particular atom of neon has a mass number of 20. How many
neutrons does is have?
Mass# = 20
Atomic# = 10 (has 10 protons)
Mass# - Atomic# = 20 – 10 = 10 Neutrons
- Sample Problem A: How many protons, neutrons, and electrons are present in
an atom of copper whose atomic number is 29 and whose mass number is 64.
Copper (Cu)
?p+ = 29
?e- = 29
?n0 = mass – atomic = 64 – 29 = 35
- The atomic number always appears on the lower left side of the symbol.
- Mass numbers are written on the upper left side of the symbol.
1
2
2
2
3
7
5
9
11
1H
1H
3He
4He
6Li
3Li
4Be
5B
5B
+
+
+
+
+
+
+
+
1p 2p 2p 2p
3p 7p 5p 9p 11p+
0n0 1n0 0n0 0n0 0n0 4n0 1n0 4n0 6n0
1e- 1e- 3e- 4e- 6e- 3e- 4e- 5e- 5e10/22/12
Atomic Number and Mass Number
Isotopes
- All atom of an element have the same atomic number and the same number of
protons. Atoms do not necessarily have the same number of neutrons.
- Atoms of the same element that have different numbers of neutrons.
- One standard method of identifying isotopes is to write the mass number with
a hyphen after the name of the element.
o helium-3 or helium-4
- The second method of identifying isotopes shoes the composition of a nucleus
as the isotope’s nuclear symbol.
o 32He or 42He
10/23/12
Do Now
-
Element
Magnesium
Magnesium
Symbol
Mg
Mg
Magnesium
chloride
Magnesium
Potassium
ClMg
KMg
Magnesium
Phosphorus
Magnesium
Aluminum
P Mg
AlMg
Magnesium
U
Mg
Au
Uranium
Magnesium
Gold
Molybdenum Mo
Cesium
Cs
p+
n0
e-
Mass#
12
17
19
15
13
92Mg
79
42
55
12
18
21
16
14
146
118
54
78
12
17
19
15
13
92
79
42
55
24
35
40
31
27
238
197
96
133
Atomic#
12
17
19
15
13
92
79
42
55
10/24/12
3.3 Electron Configuration- Atomic Models
Rutherford’s model proposed electron orbits
- The experiments of Rutherford’s team led to the replacement of the plum
pudding model of the atom with a nuclear model of the atom.
o Rutherford suggested that electrons, like planets orbiting the sun,
revolve around the nucleus in circular or elliptical orbits.
o Rutherford’s model could not explain why electrons did not crash into
the nucleus.
- Neils Bohr replaced the Rutherford model of an atom 2 years later.
Bohr’s Model Confines Electrons to Energy Levels
- According to Bohr’s model, electrons can be only certain distances from the
nucleus. Each distance corresponds to a certain quantity of energy that an
electron can have.
o An electron that is as close to the nucleus as it can be is in its lowest
energy level.
o The farther an electron is from the nucleus, the higher the energy level
that the electron occupies.
- The difference in energy between two energy levels is known as a Quantum of
energy
- Rutherford had a planetary model
- Bohr had the quantum energy model
Electrons Act like Both Particles and Waves
- Thomson’s experiments demonstrated that electrons act like particles that
have mass.
- In 1924, Louis de Broglie pointed out that the behavior of electrons according
to Bohr’s model was similar to the behavior of waves.
- De Broglie suggested that electrons could be considered waves confined to the
space around the nucleus.
- The present day model of the atom takes into account both the particle and
wave properties of electrons.
- In this model, electrons are located in Orbitals, regions around a nucleus that
correspond to specific energy levels
o Orbitals are regions where electrons are likely to be found.
10/26/12
Spectroscopy- study of light absorption, separating light into its
component colors.
Electrons and light
- By 1900, scientists knew that light could be though of as moving waves that
have given frequencies, speeds, and wavelengths.
- The wavelength is the distance between two consecutive peaks pr troughs of a
wave.
- The electromagnetic spectrum is all the frequencies or wavelengths of
electromagnetic radiation.
- In 1905, Albert Einstein proposed that light also has come properties of
particles.
His theory explained a phenomenon known as the Photoelectic Effect.
o This effect happens when light strikes a metal and electrons are
released.
- Einstein proposed that light has the properties of both particles and waves.
Light is an electromagnetic wave
- Red light has a low frequency and a long wavelength
- Violet light has a high frequency and short wavelength
-
11/6/12
Wavelength and Frequency
- The distance between any two corresponding points on one wave, such as
from crest to crest, is the wavelength
- We can measure the wave’s frequency, v, by observing how often the wave
rises and falls at a given point.
Light provides information about Electrons
- An electron in a state of its lowest possible energy is in a Ground State.
o The ground state is the lowest energy state of a quantized system.
- If an electron gains energy, it moves to an Excited State.
o An excited state is a state in which an atom has more energy than it
does at its ground state.
- An electron is an excited state will release a specific quantity of energy as it
quickly “falls” back to its ground state.
Quantum Numbers
- The present-day model of the atom is known as the Quantum Model
- A Quantum Number is a number that specified the properties of elements.
11/7/12
- The principal quantum number, symbolized by n, indicates the main energy
level occupied by the electron.
o Values of n are positive integers, such as 1, 2, 3, and 4.
o As n increases, the electron’s distance from the nucleus and the
electron’s energy increases.
- The main energy levels can be divided into sublevels. These sublevels are
represented by the angular momentum quantum number, I.
o This quantum number indicates shape or type of orbital that
corresponds to a particular sublevel.
o A letter code is used for this quantum number.
 I = 0 corresponds to an s orbital
 I = 1 to p orbital
 I = 2 to d orbital
 I = 3 to f orbital
- The magnetic quantum number, symbolized by m, is a subset of the I quantum
number.
o It also indicates the numbers and orientation of orbitals around the
nucleus.
o The value of m takes whole-number values, depending on the value of
I.
o The number of orbitals includes
 One s orbital
 Three p orbitals ( dumbbells)
 Five d orbitals
 Seven f orbitals
- The spin quantum number indicates the orientation of an electron’s magnetic
field relative to an outside magnetic field
o The spin quantum number is represented by
 + 1/2 or – 1/2 or  or 
o A single orbital can hold a maximum of two electrons, which must
have opposite spins.
11/9/12
Do Now
What are the 4 orbitals?
S p d f
How many shapes do they each have?
1 3 5 7
How many (max) e- can each orbital hold? 2 6 10 14
Electron Configuration
- Write the electron configuration for Nitrogen.
7 e1s22s22p3
14 Write the electron configuration for Calcium
20 e1s22s22p63s23p64s2
15 Write the electron configuration for Silicon
14 e
1s22s22p63s23p2
11/12/12
Do Now- Write the electron configuration for:
Mg- 12 electrons
valence electrons

2 6 2
1s22s 2p 3s

outermost shell
6 2 5
(2 valence electrons)
Cl- 1s22s22p 3s 3p
C- 1s22s22p2
He- 1s2
Ne- 1s22s22p6
The max number of e-‘s in an outer energy level is 8
11/13/12
Aufbau Principle
Lowest 
Energy
Level
Highest
Energy
Level 
11/15/12
Patterns in Element Properties
- The elements lithium, sodium, potassium, rubidium, and cesium can combine
with chlorine in a 1:1 ratio
o They form LiCl, NaCl, KCl, RbCl, and CsCl
John Newlands
- In 1865, the English chemist John Newlands arranged the known elements
according to their properties and in order of increasing atomic mass. He
placed the elements in a table.
- Newlands noticed that all of the elements in a given row had similar chemical
and physical properties
- Because these properties seemed to repeat every eight elements, Newlands
called this pattern the Law of Octaves.
Dmitri Mendeleev- invented the first periodic table
- In 1869, the Russian chemist Dmitri Mendeleev used Newlands’ observation
and other information to produce the first orderly arrangement, or periodic
table, of all 63 elements known at the time.
- Mendeleev wrote the symbol for each element, along with the physical and
chemical properties and the atomic mass and number on a card.
- Mendeleev started a new row each time he noticed the chemical properties of
the elements repeated.
o He placed elements in the new row directly below elements of similar
chemical properties in the preceding row.
- Mendeleev made two interesting observations
o Mendeleev’s table contains gaps that elements with particular
properties should fill.
o The elements do not always fit neatly in order of atomic mass.
- Mendeleev predicted the properties of the missing elements.
The Physical Benefits of the Periodic Table
- About 40 years after Mendeleev published his periodic table, an English
chemist named Henry Moseley found a different physical basis for the
arrangement of elements.
Moseley studied the lines in the X-ray spectra of 38 different elements; he
found that the wavelengths of the lines in the spectra decreased in a regular
manner as atomic mass increased.
11/16/12
- When elements were arranged by increasing atomic number, the discrepancies
of Mendeleev’s table disappeared.
Periodic Law
Organization of the Periodic Table
- When the elements are arranged by increasing atomic number, elements with
similar properties appear at regular intervals.
- Elements in each column have the same number of electrons in their outer
shell.
- The electrons in the outer shell are called Valence Electrons.
o Valence Electrons are found in the outermost shell of an atom and that
determines the atom’s chemical properties
- Elements with the same number of valence electrons tend to react in similar
ways.
- A vertical column on the periodic table is called a group. Elements in a group
share chemical properties.
- A horizontal row on the periodic table is called a period. Elements in the same
period have the same number of occupied energy levels.
11/19/12
Tour of the Periodic Table
Main Group Elements
- Elements in groups 1,2 and 13-18 are known as the main-group elements.
Main- group elements are in s- and p-blocks of the periodic table.
- Main- group elements are sometimes called representative elements because
they have a wide range of properties
o At room temperature and atmospheric pressure, many are solids, while
others are liquids or gases.
o About half of the main-group elements are metals
o Many are extremely reactive, while several are nonreactive.
- The main-group elements silicon and oxygen account for four of every five
atoms found on or near the earth’s surface.
- Four groups within the main-group elements have special names. These are:
o Alkali metals (group 1)
o Alkaline-earth metals (Group 2)
o Halogens (Group 7 or 17)
o Noble gases (Group 8 or 18)
The Alkali Metals make up Group 1
- Elements in Group 1 are called alkali metals
o Litium, sodium, potassium, rubidium, cesium, and francium
- Alkali metals are so named because they are metals that react with water to
make alkaline solutions.
- Because the alkali metals have a single valence electron, they are very
reactive.
-
Alkali metals are never found in nature as pure elements but are found as
compounds.
The Alkaline-Earth Metals Make up Group 2
- Group 2 elements are slightly less reactive than the alkali metals
- The alkaline-earth metals are slightly less reactive than the alkali metals
o They are usually found as compounds.
- The alkaline-earth metals have two valence electrons and must lose both their
valence electrons to get to a stable configuration.
o It takes more energy to lose two electrons than it takes to lose just on
electron
The Halogens make up Group 7 (or 17)
- They are the most reactive group of nonmetal elements.
o When Halogens react, they often gain the one electron needed to have
eight valence electrons, a filled outer energy level.
- Because the alkali metals have one valence electron, they are ideally suited to
react with halogens.
- The halogens react with most metals to produce salts.
The Noble Gases, Group 8 (or 18) are unreactive
- Group 8 elements are called the noble gases
- The noble gas atoms have a full set of electrons in their outermost energy
level
- The low reactivity of noble gases leads to some special uses.
- The noble gases were once called inert gases because they were thought to be
completely unreactive.
o In 1962, chemists were able to get xenon to react, making the
compound XePtF6
o In 1979, chemists were able to form the first xenon-carbon bonds.
Hydrogen is in a class by itself
- Hydrogen is the most common element in the universe.
11/27/12
Do Now
Group 1- Alkali Metals
Group 2- Alkaline Earth Metals
Group 7/17- Halogens
Group 8/18- Noble Gases
Where are the metals located? Left of the staircase
Where are the nonmetals located? Right of the staircase
How are elements organized on the periodic table? atomic number (p+)
Chapter 4 (cont)
Most elements are Metals
Transition Metals occupy the center of the periodic table
- Generally, the transition metals are less reactive than other metals.
Lanthanides and Actinides Fill f-orbitals
- The elements in the first row are called Lanthanides because their atomic
number follow the element lanthanum.
-
A lanthanide is a member of the rare-earth series of elements, whose atomic
numbers range from 58-71.
- The actinides are unique in that their nuclear structure
Other Properties of Metals
- An alloy is a solid or liquid mixture of two or more metals.
- The properties of an alloy are different from the properties of the individual
elements.
- A common alloy is brass, a mixture of copper and zinc
o Brass is harder than copper and more resistant to corrosion.
Section 3- Trends in the Periodic Table
Periodic Trends
- The arrangement of the periodic table reveals trends in the properties of the
elements.
- A trend is a predictable change in a particular direction.
- Understanding a trend among elements enables you to make predictions about
the chemical behavior of the elements.
- These trends in properties of the elements in a group or period can be
explained in terms of electron configuration.
Ionization Energy
- The ionization energy required removing an electron from an atom or ion.
A (neutral atom) + ionization energy  A+ + e11/28/12
Do Now
Define- Ductile- draw into a wire, stretch it
Malleable- bendable, folded
What “block” are the transition metals? D Block
What group are the Halogens? 7/17
Alkaline Earth Metals? 2
Noble gases? 8/18
Alkali Metals? 1
Periodic Trends
1) Atomic Radius- distance from the nucleus to the edge of the electron cloud


increases decreases
-
Electron cloud
2)
3)
4)
5)

-Bond Radius – half the distance between bonding atoms
(nucleii)
Electronegativity- the measure of the ability of an atom in a compound to attract
electrons


decreases increases
Ionic Size- the increase or decrease in the size of an atom when it gains or loses
electrons (ion- an atom with a charge)
Ionization Energy- the energy required to remove an electron from an atom or ion
BP/MP(boiling point/melting point)- changes in state due to energy changes
11/29/12
Do Now
- List the 4 evidences of a chemical change: color change, evolution of a gas,
release or absorption of energy, precipitate.
- Convert 3.70kg to grams: 3700g
- Place 23700 into scientific notation: 2.37x104
- What is the metric system unit of volume: L
12/3/12
CHAPTER 7
The Mole
- 1 mole = 6.02x1023
- Mole - defined in terms of carbon
- the number of atoms in exactly 12 grams of the carbon isotope-12
- 1 mole = atomic mass on periodic table
- Avogadro’s Number- the number of particles in 1 mole of a substance
12/4/12
Avogadro’s Number and the Mole
- The mole is a counting unit
- Amount in Moles can be converted to number of particles.
- Counting units are used to make conversion factors
- The conversion factor is 6.02x1023particles / 1mol
12/5/12
Converting Amount in Moles to Number of Particles
- Law of Conservation of Mass- xg in reactants = xg in products
- Example: Find the number if molecules in 2.5mol of sulfur dioxide?
Given: 2.5mol sulfur dioxide
? # of molecules
2.5mol sulfur dioxide
6.02x1023 molecules =
1 mol sulfur dioxide
12/6/12
MUST KNOW THIS:
1mol
or
6.02x1023
1mol
molar mass or atomic mass
or
6.02x1023
1mol
molar mass or atomic mass
1 mol
HW Review
6c) Given: 0.25molK+
? molecules
0.25molK+ 6.02x1023moleculesK+ = 1.505x1023moleculesK+
1molK+
7b) Given: 3.00molNa4P2O7
? # Na+ ions
3.00molNa4P2O7 6.02x1023Na4P2O7
4Na+ = 7.224x1024Na+
1molNa4P2O7
1Na4P2O7
8a) Given: 3.01x1023moleculesH2O
? # moles
3.01x1023moleculesH2O
1molH2O
= .5molH2O
6.02x1023moleculesH2O
12/7/12
Do Now
6.7molH2O 6.02x1023H2O = 4.0x1024moleculesH2O
1molH2O
2.68molCO2
6.02x1023CO2 = 1.61x1024moleculesCO2
1molCO2
0.005molH2O 6.02x1023H2O = 3.01x1021moleculesH2O
1molH2O
Molar Mass
- the mass, in grams, of 1 mole of an element
- g/mol
- the mass of 1 mole of an element = the element’s atomic mass (from P.T.)
- ex: 1molCu = 63.55g
- 197gAu and 63.5gCu are both 1mol which both have 6.02x1023particles
- 1molH2O = 18g
12/10/12
Do Now
- How many atoms in 2.75 moles of H2O?
2.75molH2O 6.02x1023atomsH2O = 1.65x1024atomsH2O
1molH2O
- How many moles in 1.86x1024molecules of H2SO4?
1.86x1024moleculesH2SO4
1molH2O
= 3.08molH2SO4
6.02x1023moleculesH2SO4
HW Review
13a) 2.0molH2 6.02x1023moleculesH2 = 1.2x1024moleculesH2
1molH2
b) 4.01gHF 1molHF 6.02x1023moleculesHF = 1.21x1023moleculesHF
20.01gHF
1molHF
9) 4.30x1016atomsHe
1molHe
4.00gHe = 2.85x10-7gHe
6.02x1023atomsHe 1molHe
Molar Mass- The mass, in grams (g), of 1 mole of a substance, element,
compound.
Ex:
H2O = 2H = 2x1 = 2
1O = 1x16 = 16
18g/mol
12/12/12
Percent Composition
- percent, by mass, of each element in a compound
molar mass of the part
molar mass of the whole
ex: XeF4
1Xe: 1x131= 131 %Xe: Xe = 131 100 = 63.3%
4F: 4x19 =+76
XeF4 207
207
%F: F4 = 76 100 = 36.7%
XeF4 207
ex: FeO
1Fe: 1x56= 56 %Fe: Fe = 56 100 = 78%
1O: 1x16=+16
FeO
72
72
%O: O = 16 100 = 22%
FeO
72
12/13/12
Do Now
Determine the molar mass of Mg(ClO3)2
mass of each element from P.T.
Mg(ClO3)2 1Mg: 1x24= 24
2Cl: 2x35.5= 71
6O: 6x16= +96
191g/mol molar mass of Mg(ClO3)2
Determine the %composition of NH4Br
1N: 1x14= 14
%N:
N =
4H: 4x1=
4
NH4Br
1Br: 1x80= +80
98g/mol
%H: H4 =
NH4Br
%Br:
14 100 = 14.3%
98
4 100 = 4.1%
98
Br = 80  100 = 81.6%
NH4Br
98
Formula and Percentage Composition
Empirical Formula
- An empirical formula is chemical formula that shows the simplest ratio of the
relative numbers and kinds of atoms in a compound.
- An actual formula shoes the actual ratio of elements or ions in a single unit of
a compound.
- For example, the empirical formula for ammonium nitrate is NH2O, while the
actual formula is NH4NO2.
You can use percent composition for a compound to determine its empirical
formula.
o Convert the percentage of each element to grams
o Convert from g to mol using the molar mass of each element as a
conversion factor.
Determine the empirical formula of the Do Now from today: (now divide by the smallest
-
mol)
%N:
N =
NH4Br
%H: H4 =
NH4Br
%Br: Br =
NH4Br
14 100 = 14.3% 14.3gN
98
4 100 = 4.1%
4.1gH
98
80  100 = 81.6% 81.6gBr
98
1molN = 1molN = 1
14gN
1
1molH = 4.1molH = 4
1gH
1
1molBr = 1molBr =1
80gBr
1
NH4Br
Ex: Chemical analysis of a liquid shoes that it is 60.0%C, 13.4%H, and 26.6%O
by mass. Calculate the empirical formula of this substance.
%C: 60.0% 60.0gC 1molC = 5molC = 2.9 = 3
12gC
1.7
%H: 13.4% 13.4gH 1molH = 13.4molH =7.9 = 8 C3H8O
1gH
1.7
%O: 26.6% 26.6gO 1molO = 1.7molO = 1 = 1
16gO
1.7
12/14/12
Do Now
Determine the molar mass of Ca(NO3)2
Ca: 1x40 = 40
N: 2x14 = 28
O: 6x16 = + 96
164g/mol
Determine the %composition of H2SO4
H: 2x1 = 2
%H: H = 2  100 = 2%
S: 1x32 = 32
H2SO4
98
O: 4x16 = +64
%S: S = 32  100 = 32.7% (all % add to 100)
H2SO4 98
%O: O = 64  100 = 65.3%
H2SO4
98
HW Review
2) Molecular Formula: C8H18 , so Empirical Formula is C4H9
4b) ?Empirical Formula
50.1%S 50.1gS
1molS = 1.5molS = 1
32gS
1.5
SO2
49.9%O 49.9gO 1molO = 3.1molO = 2
16gO 1.5
5) Given an experimental mass of 64g/mol, what is the molecular formula for 4b?
Since the empirical mass and experimental mass are the same, the
molecular formula is the name as the empirical formula.
6) Calcium Sulfate  CaSO4 , then determine the % composition.
Molecular Formula (actual formula)- a whole number multiple of the Empirical
Formula
12/19/12
CHAPTER 8
Chemical Reaction
- A chemical reaction is the process by which one or more substances change
into one or more new substances.
- Reactants are the original substances in a chemical reaction.
- Products are the substances that are created in a chemical reaction.
Evidence of a Chemical Reaction
- Release of energy (heat, light, sound, electricity)
- Formation of gas
- Formation of a precipitate (solid)
- Change in color or odor
Symbols:
Reactants yields Products
(s)(l)(g)- physical state
(aq)- (aqueous) dissolved in water
heat or
add energy
catalyst
name or formula of catalyst
reversible reaction
Example Chemical Reaction:
- Na2O(s) + H2O(l) 1NaOH(aq)
1/3/13
Do Now
1) 2H2 + O2  2H2O
2) N2 + 3H2  2NH3
3) 6CO2 + 6H2O  C6H12O6 + 6O2
4) 2HgO  2Hg + O2
Coefficients = Moles
1/7/13
Types of Chemical Reactions
Single Displacement Rxn- one element displaces (or replaces) another in a compound
Ex: A + BC=B + AC
Ex: 2Al(s) + 3CuCl2(aq)3Cu(s) + 1AlCl3(aq)
Double Displacement Rxn- positive and negative portions of 2 compounds are
interchanged.
Ex: AB + CD  AD + CB
Ex: HCl + NaOH  HOH + NaCl
(or H2O)
Decomposition Rxn- substances break up into simpler substances when energy is applied.
Ex: ABA + B
Ex: 2H2O electricity 2H2 + O2
Synthesis Rxn- Two or more substances combine to form a new substance.
Ex: A + B  AB
Ex: C + O2 CO2
Combustion Rxn- Oxidation reaction of an organic compound in which heat/energy is
released. Every combustion has O2, CO2, and H2O.
Ex:
+ O2 CO2 + H2O + energy
Ex: C3H8 + 5O23CO2 + 4H2O + energy
1/8/12
Do Now
- What is a hypothesis? An educated guess on what you think will happen.
- What is kinetic energy? The energy of motion.
- A chemical change produces a new substance. A physical change changes the
state of a substance.
HW Review (pg 285 1-10)
1) 2 smaller compounds come together to form a larger one.
2) They are very reactive so they replace other elements.
6) For a double displacement reaction, 2 new compounds must be produced.
8) a) single displacement
b) synthesis
c) decomposition
d) double displacement
e) single displacement
f) combustion
Mole Ratio- uses coefficients
-Molesknown (or given)mole ratiomolesunknown
-Molesgiven Molesunknown Mole
Molesgiven
ratio
C6H12O6 + 6O26CO2 + 6H2O + energy
Coefficients are in green, tell us the number of moles
Ex: C6H12O6 + 6O2  6H2O + 6CO2 + Energy
1mol C6H12O6 6molO2
6molH2O
Ex: 2Al + Fe3N2  2AlN + 3Fe
Given: 6molAl
6molAl
Unknown: ?molFe
Unknown: ?molAlN
6molAl
6molCO2
3molFe = 9molFe
2molAl
2molAlN = 6molAlN
2molAl
1/9/13
Do Now
Write 2,350,000 in scientific notation: 2.35x106
Metric System unit for Mass: grams
Volume: liters
Length: meters
What are the three subatomic particles? p+, no, enucleus
electron cloud
Examples:
Br2 + Cl2  2BrCl
- 2.74molCl2
?molBrCl
-
-
239.7gCl2
779.87gBrCl
?gBrCl
2.74molCl2
239.7gCl2 1molCl2
71gCl2
2molBrCl = 5.48molBrCl
1molCl2
2molBrCl 115.5gBrCl=
1molCl2
1molBrCl
4.53x1025molecCl2
4.53x1025molecCl2 1molCl2 1molBr2 160gBr2= 1.20x104gBr2
?gBr2
6.02x1023Cl2 1molCl2 1molBr2
1/10/13
Do Now
-
Mg-25  p+=12, no=13, e- =12
Write the electron configuration for phosphorus: 1s22s22p63s23p3
Name the 4 orbitals? S P D F
e- = 2 6 10 14
shape =1 3 5 7
- What was Thompson’s Atomic Model? plum pudding model/chocolate chip
cookie model
- How many valence electrons in Be? 2eDiatomic Elements: H2, N2, O2, F2, Cl2, Br2, I2
1/11/12
Do Now
- How is the periodic table arranged? By atomic number/number of protons
- Who developed the modern periodic table? Mendeleev
- Name of horizontal rows? Periods
- Name of vertical columns? Groups
- Name of:
o Group 1- Alkali Metals
o Group 2- Alkaline Earth Metals
o Group 7- Halogens
o Group 8- Noble Gases
o D block elements- transition metals
- Name elements right of the zigzag? nonmetals
- Name elements left of the zigzag? metals
Examples:
Br2 + 5F2  2BrF5
R  P
- Diatomic Elements
- Type of reaction- synthesis
- Balancing coefficients- moles
- Molar mass- mass in grams, of 1 mole of a substance
- Avogadro’s Number- number of particles in 1 mole of a substance
Useful Ratios:
1mol/MM or MM/1mol from periodic table
1mol/6.02x1023 or 6.02x1023/1mol Avogadro’s number
THINGS WE NEED TO COVER:
Electron Dots- put one dot on each side of the element’s symbol first, then go back and
start doubling.
Mg: (2 valence electrons)
..
·S: (6 valence electrons)
˙
..
:Ar: (8 valence electrons) (octet rule- there can be, at most, 8 valence electrons)
˙˙
Ca: (2 valence electrons)
Yield
Theoretical yield- the maximum quantity of product that a reaction could make if
everything works perfectly.
Actual yield-the mass of the product actually formed.
Why? -The actual yield is less than the theoretical yield.
- Many reactants do not completely use up limiting reactants.
- Purification
- Unwanted “side reactions”
Limiting Reactant- the substance that controls the quantity of the product that can be
formed in a chemical reaction.
The limiting reactant forms the least amount of product. (runs out)
Excess Reactant- the substance that is not used up completely in a reaction.
% yield- actual yield
x100
theoretical yield
Example: Determine the limiting reactant, theoretical yield and %yield if 14.0gN2 are
mixed with 9gH2 and 16.1gNH3 form. N2+3H22NH3
Molar mass’s- 14g
9g

16.1g (actual yield)
14gN2 1molN2 2molNH3 17gNH3 = 17gNH3 theoretical yield
28gN2 1molN2 1molNH3
(17 is smaller then 51 so N2 is the limiting reactant)
9gH2 1molH2 2molNH3 17gNH3= 51gNH3
2gH2
3molH2 1molNH3
%yield- actual yield
= 16.1g x100 = 94.7%
theoretical yield
17g