Atomic Theory
Scientists/contributions, subatomic particles, electron configuration, ions/isotopes, atomic number,
atomic mass
The atomic theory we use today was not simply discovered all at once. Over many years
scientists have made discoveries, adding to and configuring that of previous theories. Here are some of
the most important theorists.
Democritus, a Greek philosopher from the 4th century BC, theorized that if you keep cutting a substance
in half, they will eventually be too small to divide. He called these units “atoms”, after the Latin word
atomos, meaning unable to divide. Another famous Greek philosopher you’ve probably heard of,
Aristotle,
Bullets
Atomic Theory: The atomic theory we use today was not simply discovered all at once. Over many years
scientists have made discoveries, adding to and configuring that of previous theories. Here are some of
the most important theorists.
Major Contributors:
Democritus:
4th Century BC Greek philosopher
Theorized that if you keep cutting a substance in half, they will eventually be too small to divide.
He called these units “atoms”, after the Latin word atomos, meaning unable to divide
Aristotle opposed his theory
Dalton:
Interested in weather. Studied gases, concluding that they consist of individual particles
Discovered different compounds have specific ratios of element
J.J. Thompson:
Used Cathode-Ray Tube to discover electrons (+ and – metal plates attracted particles)
Made Plum-Pudding Model
Rutherford:
Discovered radioactive half-life and differentiated alpha and beta radiation
Found nuclei in gold atoms that bent alpha particles
Neil Bohr:
Danish physicist from early 20th century
Made Bohr’s model in 1911: planetary style electron cloud model we use today
Atomic Structure:
Nucleus:
Houses the protons and neutrons of the atom
Where most of the mass is
Electron Orbitals:
Electron clouds/energy levels surround the nucleus
Elements with similar electron configuration have similar qualities
Suborbitals are more specific – 4 orbitals: S – 2 houses (holds 2e-) P – 3 houses (holds 6 e-) D – 5 houses
(holds 10 e-) and F – 7 houses (holds 14e-)
Houses are “clumps” within orbitals that have to do with behavior
Protons:
Positively charged subatomic particles in nucleus
~2000x larger than electrons
Mass of 1.674x10-24
Neutrons:
Neutral subatomic particles in nucleus
~The same mass as protons (1.675x10-24
Hold protons together with strong force
Electrons:
From Greek word “elektron” meaning amber. Amber has an electrical charge
Circle atom’s nucleus in orbitals
Much smaller than protons and neutrons (9.11x10-28) and not counted in mass #
Attracted to protons by electromagnetic force
Periodic Table Design:
Atomic Number:
Amount of protons in an atom
Determines what element it is
Atomic Symbol:
Letter(s) used to show/abbreviate a certain element in chemical equations and on the periodic table
Ex: Sodium = Na; Chlorine = Cl:
Na + Cl = NaCl
Element Name:
Used to identify the element on the periodic table
Many periodic tables do not use element names, so you must memorize the symbols
Atomic Mass:
Determined by #of protons + # of neutrons
Measured in amu – atomic mass unit
Average of masses of all isotopes of an element (often weighted by the % of isotopes)
Identifying #s:
Atomic # - Number of protons (found above symbol)
Atomic Mass – p++ no (found below symbol)
Oxidation # - 1;2;3;N/A;-3;-2;-1;0 (found above each group)
Isotopes:
Atoms of the same element with different numbers of neutrons / mass #
Still retain identity an properties ( same proton # )
Written as ([Element Name] – [Mass #]) to identify
Ex: Carbon-12; Carbon-13; Carbon-14
Ion:
An atom that has gained or lost electrons to have a -/+ charge
Cations – Atoms that lose electrons to become positively charged
Anions – Atoms that gain electrons to become negatively charged
Groups:
Columns of the periodic table
Elements in the same group often have similar physical/chemical properties
Often called “families”
Periods:
Rows of the periodic table
Properties such as conductivity/reactivity change gradually down a period
Catagories:
* VE = Valence Electrons
Alkali Metals:
Metals, 1 VE, Very reactive
Soft, silver, shiny, low density
Alkaline-Earth Metals:
Metals, 2 VE, Very reactive
Silvery, greater density
Boron Group:
1 Metalloid, 5 metals, 3 VE, Reactive
Solid at room temperature, most common is aluminum
Carbon Group:
1 Nonmetal, 2 Metalloids, 3 Metals, 4 VE, Reactivity varies
Solid at room temperature
Nitrogen Group:
2 Nonmetals, 2 Metalloids, 2 Metals, 5 VE, Reactivity varies
Solid at room temp. (except nitrogen)
Oxygen Group:
3 Nonmetals, 1 Metalloid, 1 Metal, 6 VE, Reactive
Solid at room temp. (except oxygen)
Halogens:
Nonmetals, 7 VE, Very reactive
Poor conductors, violent reaction with Alkali Metals to form salts; never pure in nature
Noble Gases:
Nonmetals, stable, unreactive
Colorless, odorless gas at room temp.
Periodic Trends:
Reactivity:
A measure of how likely an element is to react with others
Decreases across a period
Radius:
The distance between an atom's nucleus and its valence electrons
Decreases across a period and increases down a group
Ionization Energy:
The amount of energy required for a chemical reaction to occur
Increases across a period and decreases down a group
Electronegativity:
Ability of an atom to attract electrons of another atom when they are associated through a bond
Based on atom's ionization energy and electron affinity
Increases across a period and decreases down a group
Chemical Bonding:
Reaction Types:
Synthesis:
In a synthesis reaction two or more atoms combine
EX: Na + Cl → NaCl
Decomposition:
A compound is broken down into simpler substances
Ex: 2H2O → 2H2 + 2O
Single Displacement:
An element in a compound is replaced by another
Ex: Zn + 2HCl → ZnCl2 + H2
Double Replacement:
Elements in two different compounds trade places with each other
Ex: NaCl + AgNO3 → NaNO3 + AgCl
Ionic bonds:
Metal + nonmetal
The metal (cation) loses electrons and becomes positively charged. The nonmetal (anion) gains those
electrons and becomes negatively charged. They are then attracted by electromagnetic force and bond.
The oxidation number of an element tells how many electrons they lose or gain in an ionic bond. The
oxidation number of transition metals varies.
Ex: Na+ClCovalent Bonds:
Nonmetal + nonmetal
Atoms share electrons in a covalent bond
In either a single (-) or double (=) bond
Ex: CH4 (methane) in which 1 carbon atom shares a single bond with 4 different hydrogen atoms.
Metallic Bonds:
Metal + metal
Sharing of electrons among positively charged metal ions (cations)
Often used in contrast to covalent bonds
Compounds / Equations
Naming:
Ionic:
Use roman numerals in parentheses next to metal to show oxidation #
(Cation)(Anion)
Ex: Ni2(SO3)3 – Nickel (III) sulfite
Covalent:
Element with lower group # first
If same, greater period first
2nd is anion (ends in ide)
Use Greek prefixes:
1 – mono
2 – di
3 – tri
4 – tetra
5 – penta
6 – hexa
7 – septa
8 – octa
9 – nona
10 – deca
How to remember: COP – Covalent prefix
Ex: P3N4 – Triphosphorus tetranitride
Naming Organic Compounds:
Alkanes:
Hydrocarbons
Hydrogen – Carbon ratio is at max, so they are saturated (formula of CnH2n+2)
1. Find and name the longest continuous carbon chain.
2. Identify and name groups attached to this chain.
Group
CH3–
C2H5–
CH3CH2CH2–
(CH3)2CH–
CH3CH2CH2CH2–
(CH3)2CHCH2–
CH3CH2CH(CH3)–
(CH3)3C–
Name
Methyl
Ethyl
Propyl
Isopropyl
Butyl
Isobutyl
sec-Butyl
tert-Butyl
R–
Alkyl
3. Number the chain consecutively, starting at the end nearest a substituent group.
4. Designate the location of each substituent group by an appropriate number and name.
5. Assemble the name, listing groups in alphabetical order using the full name (e.g. cyclopropyl before
isobutyl).
The prefixes di, tri, tetra etc., used to designate several groups of the same kind, are not considered
when alphabetizing
Cycloalkanes:
Alkenes and Alkynes:
Benzene Derivatives:
Compounds:
Balancing Equations:
Both sides of the equation must have the same amount of each element
You cannot change subscripts, only coefficients, to balance equations
Ex: NaCl  2Na + 2Cl
You must add a coefficient of 2 in front of NaCl
When the equations get more complicated it’s easier to use the criss-cross method. List the elements
and how many there are of each on both sides of the equation. Then match up which elements are
inconsistent. Choose one and fix it. Repeat the process until the equation is balanced.
Sources:
http://www.sparknotes.com/chemistry/
http://chemistry.about.com/od/chemicalreactions/a/reactiontypes.htm
https://www2.chemistry.msu.edu/faculty/reusch/virttxtjml/nomen1.htm
Along with some of my own sources!
Photos:
http://msp.ehe.osu.edu/wiki/images/d/d7/Tp_atomic.jpg
http://www.dvc.edu/org/images/beakers.jpg
http://rightbrainaerobics.com/IMAGES/A152120-Atomic_structure-SPL.jpg
http://farm3.static.flickr.com/2629/4004302303_c58c634622_z.jpg
http://graphiceducationonline.com/media/catalog/product/cache/1/thumbnail/600x600/9df78eab3352
5d08d6e5fb8d27136e95/g/e/general_periodic_trends_1.jpg
http://staff.jccc.net/pdecell/chemistry/water.gif
http://www.chem.qmul.ac.uk/iupac/home/IUPAC.GIF
http://2012books.lardbucket.org/books/introduction-to-chemistry-general-organic-andbiological/section_15/445d51e87ee3d8315d8f95e17a01e597.jpg