Name __________________________________ Period _________ Date _________________
Chemistry Unit #1 – Atomic Concepts Study Guide

The modern model of the atom has evolved over a long period of time through the work
of many scientists.
1. Plum Pudding Model – composed of a positive charge (pudding) with some
random electrons (plum pieces)
2. Rutherford Model – Gold Foil Experiment (positive center, surrounded by
scattered electrons)
3. Bohr Model – Positive center with neutron, surround by electrons in an specific
orbit
1. Know how to draw Bohr models with appropriate # of protons and
neutrons at the center of the atom
2. Proper placement of electrons around the nucleus
3. Remember that there can only be a max of 2 electrons on the 1st shell,
and a max of on the 2nd shell.
4. Wave Mechanical Model – Electrons are found in wave like orbitals at specific
energy levels – no exact path, probability to estimate the location of electron.
1. S- orbital (can hold a max of 2 electrons)
2. P – orbital ( can hold a max of 6 electrons)
3. D- orbital (can hold a max of 10 electrons)
4. F – orbital (can hold a max of 14 electrons)

Each atom has a nucleus, with an overall positive charge, surrounded by negatively (-)
charged electrons.

Subatomic particles contained in the nucleus include protons and neutrons.

The proton is positively (+) charged, and the neutron has no charge. The electron is
negatively charged (-).

Protons (+1) and electrons (-1) have equal but opposite charges. The number of protons
equals the number of electrons in an atom.

Mass of each proton and each neutron is approximately equal to one atomic mass unit.
An electron is much less massive than a proton or a neutron.

Mass = Proton + neutron; to find neutron = mass - proton

In the wave-mechanical model (electron cloud model) the electrons are in orbitals, which
are defined as the regions of the most probable electron location (ground state).
1. Ground state to excited state (ex. 2-8-1  2-7-2) Energy was absorbed by
electron.
2. Excited state to ground state (ex. 2-7-2  2-8-1) Energy was emitted/released by
electron.
3. Make sure you know how to draw for both Bohr Models.

Each electron in an atom has its own distinct amount of energy.

When an electron in an atom gains a specific amount of energy, the electron is at a higher
energy state (excited state).

When an electron returns from a higher energy state to a lower energy state, a specific
amount of energy is emitted. This emitted energy can be used to identify an element.

The outermost electrons in an atom are called the valence electrons. In general, the
number of valence electrons affects the chemical properties of an element.
1. (As you go down the period, elements have similar chemical properties) (ex. O, S,
Se, Te have similar properties)

Atoms of an element that contain the same number of protons but a different number of
neutrons are called isotopes of that element.
2. (Ex. Carbon-14, carbon – 13, carbon -12) Different mass, same # of protons,
different number of neutrons)

The average atomic mass of an element is the weighted average of the masses of its
naturally occurring isotopes.
3. Average atomic mass = (% abundance x mass of 1st isotope) + (% abundance x
mass of 2st isotope) + …
(Remember to convert the percentage into a decimal first!)

Lewis Dot Diagrams
4. Make sure to write the element symbol with its corresponding valence electrons
around it. (Hint: look at the last # on its electron configuration)

Electron Configuration
 2+ 2 + 6 + 2 = 12 electrons .:. 12 protons .:. Atomic # = 12
The element above is Magnesium
- Identify the sublevel notation (like above)
- Know how to write a sublevel notation
-
Abbreviated notation
Unabbreviated notation
Orbital notation
Recall the following…
Unit #2 Exam Study Guide – The Periodic Table
Ms. Pacheco
History of the Periodic Table

Dmitri Mendeleev: Chemist who helped develop the periodic table with its arrangement
based on the atomic number of an element as opposed to its atomic mass. (BOLD #)
 Periodic Law: elements are arranged in order of increasing atomic number (1H, 2He,
3Li…)
 Elements cannot be broken down by chemical change.
 Elements are also grouped in columns due to an element’s similar chemical properties
(Ex. Oxygen, Sulfur, Selenium, etc)
 Groups: #1-18 (in columns)
1. Group 1  Alkali metals (very reactive, most reactive METAL, not found in
nature and kept in oil, 1 valence electron)
2. Group 2  Alkaline Earth metals (reactive, 2 valence electrons)
3. Group 3 – 12  Transition metals (semi-reactive, make colored ion solutions)
4. Group 13  (3 valence electrons)
5. Group 14  (4 valence electrons)
6. Group 15  (5 valence electrons)
7. Group 16  (6 valence electrons)
8. Group 17  Halogens (7 valence electrons, most reactive NONMETAL, F & Cl
are gases, Br – liquid, I – purple solid))
9. Group 18  Noble Gas (8 valence electrons, most stable gas, not reactive,
odorless, colorless)
Properties of Elements
 Metals:
1. Solids at room temperature
2. Most have densities greater than water; except G1
3. Malleable – shaped or shaped easily
4. Ductile – stretched or pulled into a thin wire
5. Luster – shiny
6. Good conductors of heat & electricity
7. Lose electrons
 Nonmetals:
1. Gases at room temp; Br is a liquid at room temp.
2. Not malleable
3. Not ductile
4. Lackluster; appears dull
5. Poor conductors of heat and electricity
6. Gain electrons


Metalloids:
1. B, Si, Ge, As, Sb, Se and Te
2. Both have metallic and non metallic properties
3. Semiconductors
Elements can be differentiated by physical properties. Physical properties of substances,
such as density, conductivity, malleability, solubility, and hardness, differ among
elements.
Elements can also be differentiated by chemical properties. Chemical properties describe how an element behaves
during a chemical reaction


Group Trends
1. Ionization Energy: Increases as you go up (bottom’s up)
a. Highest Ionization Energy (He) found in Table S
2. Electronegativity: Increases as you go up (bottom’s up)
3. Atomic Radius: Decreases as you go up
4. Metallic Properties: Decreases as you go up
Period Trends
1. Ionization Energy: Increases from left to right
2. Electronegativity: Increases from left to right
3. Atomic Radius: Decreases from left to right
4. Metallic Properties: Decreases left to right (G1 – metals, G18 – nonmetals)
Ionization Energy: The process that results in the formation of an ion
Ion: An atom that gains/molecules that has gained or lost one or more electrons and has a (-) or
(+) charge
Positive Ion/Cation: An atom who loses electrons (+)
Negative Ion/Anion: An atom who gains electrons (-)
Remember if the atom has a (+) sign then it has lost an electron(s) thus making the atom more
positive.
Remember if the atom has a (-) sign then it has gained an electron(s) thus making the atom more
negative.
The Octet Rule: atoms gain or lose electrons in order to attain an electron configuration of the
nearest noble gas(Group 18)
Ex. Na has only 1e, thus, will most likely lose it to become its nearest noble gas which is Ne,
Neon.
2-8-1  2-8
Cl has 7e, and this is more likely to gain an electron to become its nearest noble gas which is Ar.
2-8-7  2-8-8
Types of Chemical Bonds
Chemical Bond: a strong force of attraction that keeps atoms together
Ionic Bonding: the transfer of valence electrons between atoms
Metals and nonmetals:
Covalent Bonding: the sharing of valence electrons between atoms
Nonmetals and nonmetals:
Forms single bonds, double and triple bonds
Metallic Bonding: attraction between valence electrons & metal ions.
Compound: substance in which two or more different elements are chemically combined
(# of atoms X # of valence electrons) + (# of atoms X # of valence electrons) + …
CO2
 (2C x 4ve) + (2O x 6ve) = 8+12 = 20 valence electrons on a moleciule
Polar covalent Bonds: Unequal sharing of electrons in a bond
Nonpolar covalent bonds: Equal sharing of electrons in a bond
Electronegativity
Difference
0-0.4
0.4 < n <1.7
1.7, 1.7<
Bond Type
Non-Polar Covalent
Polar
Covalent
Ionic
Example
O2
H-Cl
NaCl
Remember that even though a electronegative difference may be 1.7 or more but its between two
nonmetals, then it is considered polar covalent. (Ex. HF)
Naming Binary Compounds
Metals (fixed oxidation) + nonmetals  CsF  Cesium Fluoride
**Keep the name of the first metal element and change the ending of the second nonmetal to ide
Criss-Cross Rule