Ms. Roman’s Amazing Chemistry
Powerpoint Presentation!
 Aligned to the New York State Standards and
Core Curriculum for “The Physical SettingChemistry”
 Can be used in any high-school chemistry class!
 Please give the link to this file to your chemistry
students! www.markrosengarten.com
 Enjoy it!!! A LOT of work has gone into
bringing you this work, so please credit me when
you use it!
Outline for Review
1) The Atom (Nuclear, Electron Config)
2) Matter (Phases, Types, Changes)
3) Bonding (Periodic Table, Ionic, Covalent)
4) Compounds (Formulas, Reactions, IMAF’s)
5) Math of Chemistry (Formula Mass, Gas Laws,
Neutralization, etc.)
6) Kinetics and Thermodynamics (PE Diagrams, etc.)
7) Acids and Bases (pH, formulas, indicators, etc.)
8) Oxidation and Reduction (Half Reactions, Cells, etc.)
9) Organic Chemistry (Hydrocarbons, Families, Reactions)
The Atom
1) Nucleons
2) Isotopes
3) Natural Radioactivity
4) Half-Life
5) Nuclear Power
6) Electron Configuation
7) Development of the Atomic Model
Nucleons
 Protons: +1 each, determines identity of element, mass of





1 amu, determined using atomic number, nuclear charge
Neutrons: no charge, determines identity of isotope of an
element, 1 amu, determined using mass number - atomic
number (amu = atomic mass unit)
32 S and 33 S are both isotopes of S
16
16
S-32 has 16 protons and 16 neutrons
S-33 has 16 protons and 17 neutrons
All atoms of S have a nuclear charge of +16 due to the 16
protons.
Isotopes
 Atoms of the same element MUST contain the same
number of protons.
 Atoms of the same element can vary in their numbers of
neutrons, therefore many different atomic masses can exist
for any one element. These are called isotopes.
 The atomic mass on the Periodic Table is the weightaverage atomic mass, taking into account the different
isotope masses and their relative abundance.
 Rounding off the atomic mass on the Periodic Table will
tell you what the most common isotope of that element is.
Weight-Average Atomic Mass
 WAM = ((% A of A/100) X Mass of A) + ((% A of B/100) X Mass of B) + …
 What is the WAM of an element if its isotope masses and
abundances are:
– X-200: Mass = 200.0 amu, % abundance = 20.0 %
– X-204: Mass = 204.0 amu, % abundance = 80.0%
– amu = atomic mass unit (1.66 × 10-27 kilograms/amu)
Most Common Isotope
 The weight-average atomic mass of Zinc is
65.39 amu. What is the most common isotope
of Zinc? Zn-65!
 What are the most common isotopes of:
– Co
–S
Ag
Pb
 FACT: one atomic mass unit (1.66 × 10-27
kilograms) is defined as 1/12 of the mass of an
atom of C-12.
 This method doesn’t always work, but it usually
does. Use it for the Regents exam.
Electron Configuration
 Basic Configuration
 Valence Electrons
 Electron-Dot (Lewis Dot) Diagrams
 Excited vs. Ground State
 What is Light?
Basic Configuration
 The number of electrons is determined from the atomic
number.
 Look up the basic configuration below the atomic number
on the periodic table. (PEL: principal energy level = shell)
 He: 2 (2 e- in the 1st PEL)
 Na: 2-8-1 (2 e- in the 1st PEL, 8 in the 2nd and 1 in the
3rd)
 Br: 2-8-18-7 (2 e- in the 1st PEL, 8 in the 2nd, 18 in the
3rd and 7 in the 4th)
Valence Electrons
 The valence electrons are responsible for all chemical




bonding.
The valence electrons are the electrons in the outermost
PEL (shell).
He: 2 (2 valence electrons)
Na: 2-8-1 (1 valence electron)
Br: 2-8-18-7 (7 valence electrons)
 The maximum number of valence electrons an atom can
have is EIGHT, called a STABLE OCTET.
Electron-Dot Diagrams
 The number of dots equals the number of valence
electrons.
 The number of unpaired valence electrons in a nonmetal
tells you how many covalent bonds that atom can form
with other nonmetals or how many electrons it wants to
gain from metals to form an ion.
 The number of valence electrons in a metal tells you how
many electrons the metal will lose to nonmetals to form an
ion. Caution: May not work with transition metals.
 EXAMPLE DOT DIAGRAMS
Example Dot Diagrams
Carbon can also have this dot diagram, which it
has when it forms organic compounds.
Excited vs. Ground State
 Configurations on the Periodic Table are ground state
configurations.
 If electrons are given energy, they rise to higher energy
levels (excited state).
 If the total number of electrons matches in the
configuration, but the configuration doesn’t match, the
atom is in the excited state.
 Na (ground, on table): 2-8-1
 Example of excited states: 2-7-2, 2-8-0-1, 2-6-3
What Is Light?
 Light is formed when electrons drop from
the excited state to the ground state.
 The lines on a bright-line spectrum come
from specific energy level drops and are
unique to each element.
 EXAMPLE SPECTRUM
EXAMPLE SPECTRUM
This is the bright-line spectrum of hydrogen. The top
numbers represent the PEL (shell) change that produces the
light with that color and the bottom number is the
wavelength of the light (in nanometers, or 10-9 m).
No other element has the same bright-line spectrum as
hydrogen, so these spectra can be used to identify
elements or mixtures of elements.
Development of the Atomic
Model
 Thompson Model
 Rutherford Gold Foil Experiment and
Model
 Bohr Model
 Quantum-Mechanical Model
Thompson Model
 The atom is a positively charged diffuse
mass with negatively charged electrons
stuck in it.
Rutherford Model
 The atom is made of a small, dense, positively charged
nucleus with electrons at a distance, the vast majority of
the volume of the atom is empty space.
Alpha particles shot
at a thin sheet of gold
foil: most go through
(empty space). Some
deflect or bounce off
(small + charged
nucleus).
Bohr Model
 Electrons orbit around the nucleus in energy levels (shells).
Atomic bright-line spectra was the clue.
Quantum-Mechanical Model
 Electron energy levels are wave functions.
 Electrons are found in orbitals, regions of space where an
electron is most likely to be found.
 You can’t know both where the electron is and where it is
going at the same time.
 Electrons buzz around the nucleus like gnats buzzing
around your head.
Matter
1) Properties of Phases
2) Types of Matter
3) Phase Changes
Properties of Phases
 Solids: Crystal lattice (regular geometric pattern),
vibration motion only
 Liquids: particles flow past each other but are still
attracted to each other.
 Gases: particles are small and far apart, they travel in a
straight line until they hit something, they bounce off
without losing any energy, they are so far apart from each
other that they have effectively no attractive forces and
their speed is directly proportional to the Kelvin
temperature (Kinetic-Molecular Theory, Ideal Gas Theory)
Solids
The positive and
negative ions
alternate in the
ionic crystal lattice
of NaCl.
Liquids
When heated, the ions move
faster and eventually
separate from each other to
form a liquid. The ions are
loosely held together by the
oppositely charged ions, but
the ions are moving too fast
for the crystal lattice to stay
together.
Gases
Since all gas molecules spread out
the same way, equal volumes of
gas under equal conditions of
temperature and pressure will
contain equal numbers of
molecules of gas. 22.4 L of any
gas at STP (1.00 atm and 273K)
will contain one mole
(6.02 X 1023) gas molecules.
Since there is space between gas
molecules, gases are affected by
changes in pressure.
Types of Matter
 Substances (Homogeneous)
– Elements (cannot be decomposed by chemical
change): Al, Ne, O, Br, H
– Compounds (can be decomposed by chemical
change): NaCl, Cu(ClO3)2, KBr, H2O, C2H6
 Mixtures
– Homogeneous: Solutions (solvent + solute)
– Heterogeneous: soil, Italian dressing, etc.
Elements
 A sample of lead atoms (Pb). All
atoms in the sample consist of lead,
so the substance is homogeneous.
 A sample of chlorine atoms (Cl).
All atoms in the sample consist of
chlorine, so the substance is
homogeneous.
Compounds
 Lead has two charges listed, +2
and +4. This is a sample of lead
(II) chloride (PbCl2). Two or
more elements bonded in a
whole-number ratio is a
COMPOUND.
 This compound is formed from
the +4 version of lead. This is
lead (IV) chloride (PbCl4).
Notice how both samples of lead
compounds have consistent
composition throughout?
Compounds are homogeneous!
Mixtures
 A mixture of lead atoms and
chlorine atoms. They exist in no
particular ratio and are not
chemically combined with each
other. They can be separated by
physical means.
 A mixture of PbCl2 and PbCl4
formula units. Again, they are in
no particular ratio to each other
and can be separated without
chemical change.
Bonding
1) The Periodic Table
2) Ions
3) Ionic Bonding
4) Covalent Bonding
5) Metallic Bonding
The Periodic Table
 Metals
 Nonmetals
 Metalloids
 Chemistry of Groups
 Electronegativity
 Ionization Energy
Metals
 Have luster, are malleable and ductile, good
conductors of heat and electricity
 Lose electrons to nonmetal atoms to form
positively charged ions in ionic bonds
 Large atomic radii compared to nonmetal atoms
 Low electronegativity and ionization energy
 Left side of the periodic table (except H)
Nonmetals
 Are dull and brittle, poor conductors
 Gain electrons from metal atoms to form
negatively charged ions in ionic bonds
 Share unpaired valence electrons with other
nonmetal atoms to form covalent bonds and
molecules
 Small atomic radii compared to metal atoms
 High electronegativity and ionization energy
 Right side of the periodic table (except Group 18)
Metalloids
 Found lying on the jagged line between metals and
nonmetals flatly touching the line (except Al and Po).
 Share properties of metals and nonmetals (Si is shiny like a
metal, brittle like a nonmetal and is a semiconductor).
Chemistry of Groups
 Group 1: Alkali Metals
 Group 2: Alkaline Earth Metals
 Groups 3-11: Transition Elements
 Group 17: Halogens
 Group 18: Noble Gases
 Diatomic Molecules
Group 1: Alkali Metals
 Most active metals, only found in compounds in




nature
React violently with water to form hydrogen gas
and a strong base: 2 Na (s) + H2O (l)  2 NaOH
(aq) + H2 (g)
1 valence electron
Form +1 ion by losing that valence electron
Form oxides like Na2O, Li2O, K2O
Group 2: Alkaline Earth Metals
 Very active metals, only found in compounds in




nature
React strongly with water to form hydrogen gas
and a base:
– Ca (s) + 2 H2O (l)  Ca(OH)2 (aq) + H2 (g)
2 valence electrons
Form +2 ion by losing those valence electrons
Form oxides like CaO, MgO, BaO
Groups 3-11: Transition Metals
 Many can form different possible charges of ions
 If there is more than one ion listed, give the charge as a
Roman numeral after the name
 Cu+1 = copper (I) Cu+2 = copper (II)
 Compounds containing these metals can be colored.
Group 17: Halogens
 Most reactive nonmetals
 React violently with metal atoms to form halide




compounds: 2 Na + Cl2  2 NaCl
Only found in compounds in nature
Have 7 valence electrons
Gain 1 valence electron from a metal to form -1
ions
Share 1 valence electron with another nonmetal
atom to form one covalent bond.
Group 18: Noble Gases
 Are completely nonreactive since they have eight
valence electrons, making a stable octet.
 Kr and Xe can be forced, in the laboratory, to give
up some valence electrons to react with fluorine.
 Since noble gases do not naturally bond to any
other elements, one atom of noble gas is
considered to be a molecule of noble gas. This is
called a monatomic molecule. Ne represents an
atom of Ne and a molecule of Ne.
Diatomic Molecules
 Br, I, N, Cl, H, O and F are so reactive that they exist in a
more chemically stable state when they covalently bond
with another atom of their own element to make two-atom,
or diatomic molecules.
 Br2, I2, N2, Cl2, H2, O2 and F2
 The decomposition of water: 2 H2O  2 H2 + O2
Electronegativity
 An atom’s attraction to electrons in a chemical bond.
 F has the highest, at 4.0
 Fr has the lowest, at 0.7
 If two atoms that are different in EN (END) from each
other by 1.7 or more collide and bond (like a metal atom
and a nonmetal atom), the one with the higher
electronegativity will pull the valence electrons away from
the atom with the lower electronegativity to form a (-) ion.
The atom that was stripped of its valence electrons forms a
(+) ion.
 If the two atoms have an END of less than 1.7, they will
share their unpaired valence electrons…covalent bond!
Ionization Energy
 The energy required to remove the most loosely held
valence electron from an atom in the gas phase.
 High electronegativity means high ionization energy
because if an atom is more attracted to electrons, it will
take more energy to remove those electrons.
 Metals have low ionization energy. They lose electrons
easily to form (+) charged ions.
 Nonmetals have high ionization energy but high
electronegativity. They gain electrons easily to form (-)
charged ions when reacted with metals, or share unpaired
valence electrons with other nonmetal atoms.
Ions
 Ions are charged particles formed by the gain or loss of
electrons.
– Metals lose electrons (oxidation) to form (+) charged
cations.
– Nonmetals gain electrons (reduction) to form (-)
charged anions.
 Atoms will gain or lose electrons in such a way that they
end up with 8 valence electrons (stable octet).
– The exceptions to this are H, Li, Be and B, which are
not large enough to support 8 valence electrons. They
must be satisfied with 2 (Li, Be, B) or 0 (H).
Metal Ions (Cations)
 Na: 2-8-1
 Na+1: 2-8
 Ca: 2-8-8-2
Note that when the atom
loses its valence electron,
the next lower PEL
becomes the valence PEL.
 Ca+2: 2-8-8
 Al: 2-8-3
 Al+3: 2-8
Notice how the dot
diagrams for metal ions
lack dots! Place brackets
around the element symbol
and put the charge on the
upper right outside!
Nonmetal Ions (Anions)
 F: 2-7
 F-1: 2-8
 O: 2-6
 O-2: 2-8
 N: 2-5
 N-3: 2-8
Note how the ions all have 8
valence electrons. Also note the
gained electrons as red dots.
Nonmetal ion dot diagrams show
8 dots, with brackets around the
dot diagram and the charge of
the ion written to the upper right
side outside the brackets.
Ionic Bonding
 If two atoms that are different in EN (END) from each
other by 1.7 or more collide and bond (like a metal atom
and a nonmetal atom), the one with the higher
electronegativity will pull the valence electrons away from
the atom with the lower electronegativity to form a (-) ion.
The atom that was stripped of its valence electrons forms a
(+) ion.
 The oppositely charged ions attract to form the bond. It is
a surface bond that can be broken by melting or dissolving
in water.
 Ionic bonding forms ionic crystal lattices, not molecules.
Example of Ionic Bonding
Covalent Bonding
 If two nonmetal atoms have an END of 1.7 or less, they
will share their unpaired valence electrons to form a
covalent bond.
 A particle made of covalently bonded nonmetal atoms is
called a molecule.
 If the END is between 0 and 0.4, the sharing of electrons is
equal, so there are no charged ends. This is NONPOLAR
covalent bonding.
 If the END is between 0.5 and 1.7, the sharing of electrons
is unequal. The atom with the higher EN will be d- and the
one with the lower EN will be d+ charged. This is a
POLAR covalent bonding. (d means “partial”)
Examples of Covalent Bonding
Metallic Bonding
 Metal atoms of the same element bond with each other by
sharing valence electrons that they lose to each other.
 This is a lot like an atomic game of “hot potato”, where
metal kernals (the atom inside the valence electrons) sit in
a crystal lattice, passing valence electrons back and forth
between each other).
 Since electrons can be forced to travel in a certain direction
within the metal, metals are very good at conducting
electricity in all phases.
Compounds
1) Types of Compounds
2) Formula Writing
3) Formula Naming
4) Empirical Formulas
5) Molecular Formulas
6) Types of Chemical Reactions
7) Balancing Chemical Reactions
8) Attractive Forces
Types of Compounds
 Ionic: made of metal and nonmetal ions. Form an ionic
crystal lattice when in the solid phase. Ions separate when
melted or dissolved in water, allowing electrical
conduction. Examples: NaCl, K2O, CaBr2
 Molecular: made of nonmetal atoms bonded to form a
distinct particle called a molecule. Bonds do not break
upon melting or dissolving, so molecular substances do not
conduct electricity. EXCEPTION: Acids [H+A- (aq)]
ionize in water to form H3O+ and A-, so they do conduct.
 Network: made up of nonmetal atoms bonded in a
seemingly endless matrix of covalent bonds with no
distinguishable molecules. Very high m.p., don’t conduct.
Ionic Compounds
Ionic Crystal Structure, then adding heat (or dissolving in water) to break
up the crystal into a liquid composed of free-moving ions.
Molecular Compounds
Network Solids
Network solids are made of nonmetal atoms covalently
bonded together to form large crystal lattices. No individual
molecules can be distinguished. Examples include C
(diamond) and SiO2 (quartz). Corundum (Al2O3) also forms
these, even though Al is considered a metal. Network solids
are among the hardest materials known. They have
extremely high melting points and do not conduct electricity.
Formula Writing
 The charge of the (+) ion and the charge of the (-) ion must





cancel out to make the formula. Use subscripts to indicate
how many atoms of each element there are in the
compound, no subscript if there is only one atom of that
element.
Na+1 and Cl-1 = NaCl
Ca+2 and Br-1 = CaBr2
Al+3 and O-2 = Al2O3
Zn+2 and PO4-3 = Zn3(PO4)2
Try these problems!
Attractive Forces
 Molecules have partially charged ends. The d+ end of one
molecule attracts to the d- end of another molecule.
 Ions are charged (+) or (-). Positively charged ions attract
other to form ionic bonds, a type of attractive force.
 Since partially charged ends result in weaker attractions
than fully charged ends, ionic compounds generally have
much higher melting points than molecular compounds.
 Determining Polarity of Molecules
 Hydrogen Bond Attractions
Determining Polarity of
Molecules
-----------------------------------------------------------------------------
Hydrogen Bond
Attractions
A hydrogen bond attraction is a
very strong attractive force
between the H end of one polar
molecule and the N, O or F end
of another polar molecule. This
attraction is so strong that water
is a liquid at a temperature
where most compounds that are
much heavier than water (like
propane, C3H8) are gases. This
also gives water its surface
tension and its ability to form a
meniscus in a narrow glass tube.
Sig Figs and Rounding
 How many Significant Figures does a number have?
 What is the precision of my measurement?
 How do I round off answers to addition and subtraction
problems?
 How do I round off answers to multiplication and division
problems?
How many Sig Figs?
 Start counting sig figs at the first non-zero.
 All digits except place-holding zeroes are sig figs.
Measurement
# of Sig Figs
Measurement
# of Sig Figs
0.115 cm
3
234 cm
3
0.00034 cm
2
67000 cm
2
0.00304 cm
3
_
45000 cm
4
0.0560 cm
3
560. cm
3
0.00070700 cm
5
560.00 cm
5
What Precision?
 A number’s precision is determined by the furthest
(smallest) place the number is recorded to.
 6000 mL : thousands place
 6000. mL : ones place
 6000.0 mL : tenths place
 5.30 mL : hundredths place
 8.7 mL : tenths place
 23.740 mL : thousandths place
Rounding with addition and
subtraction
 Answers are rounded to the least precise place.
1) 4.732 cm
16.8
cm
+ 0.781 cm
---------22.313 cm
22.3 cm
2)
17.440 mL
3.895 mL
+ 16.77 mL
-------------38.105 mL
38.11 mL
3)
32.0
MW
+ 0.0059 MW
--------------32.0059 MW
32.0 MW
Rounding with multiplication
and division
 Answers are rounded to the fewest number of significant
figures.
1)
37.66 KW
x 2.2 h
---------82.852 KWh
83 KWh
2)
14.922 cm
x 2.0 cm
----------2
29.844 cm
2
30. cm
3) 98.11 kg
x 200 m
---------19 622 kgm
20 000 kgm
Metric Conversions
 Determine how many powers of ten
difference there are between the two
units (no prefix = 100) and create a
conversion factor. Multiply or divide
the given by the conversion factor.
How many kg are in 38.2 cg?
(38.2 cg) /(100000 cg/kg) = 0.000382 km
How many mL in 0.988 dL?
(0.988 dg) X (100 mL/dL) = 98.8 mL