Download Atoms, and Elements

Atoms, Molecules and Ions
I.
Atomic Theory
A. Dalton’s Postulates (1808)
1. An element is composed of tiny, indivisible particles called atoms.
2. Atoms of a given element have the same size, mass, and
properties whereas atoms of different elements have different
size, mass and properties.
3. Compounds are formed when atoms of two or more elements
combine in small whole number ratios – supports law of definite
composition and law of multiple proportion
4. In a chemical reaction, atoms move from one substance to
another but no atoms can disappear or appear from nowhere.supports Law of conservation
Part 1 and 2 of Dalton’s theory has been modified after experimental evidence
proved them incorrect
 Atoms were discovered divisible (made up of subatomic particles) between
1850-1900
 Isotopes (atoms of the same element but with different masses) were
discovered
ATOM – defined as the smallest particle of an element that can enter into a chemical
reaction
B. Basic Laws of Chemistry explained the postulates of the Atomic Theory
1. Lavoisier’s Law of conservation of mass: matter can neither be
created or destroyed but changed in form, the mass of reactants
equal the mass of products
2. Proust’s Law of definite composition(1799): a compound always
contains the same elements in the same proportions by mass
even if we look at different samples
3. Dalton’s Law of multiple proportions: different compounds made
up of the same elements; the masses of 1 element to the other
are in a ratio of small whole numbers
II.
Components of the atom
Electron Cloud
The Electron: negatively charged particles located in the outer regions of
the atom
1. JJ.Thomson discovered its charge using a cathode ray tube
2. relative charge of –1
3. Millikan discovered its actual mass
4. Actual mass = 9.109 x 10-28 g relative mass 1 amu;; approximately
2000 times smaller than the proton or neutron
 JJ Thomson formulated the PLUM PUDDING MODEL. (the positively charged
pudding had negatively charged electrons(plums) embedded in it)
Nucleus
Discovered by Rutherford in the Gold Foil experiment (1910) ;alpha particles (+ helium
nuclei) were shot through a piece of gold foil. It was assumed that the rays would come
out on the other end but some were deflected and even bounced back causing the
belief that there was a small, dense, positively charged area in the atom
The Proton
1. relative charge +1
2. relative mass 1.00728 amu or actual mass of 1.6726 x 10 -24 g
The Neutron- discovered by Chadwick(1932)
1. relative charge : neutral
1. relative mass 1.00867 amu or 1.6749 x 10-24 g
 If the Houston Astrodome was an atom, a marble placed in the stadium would be
the size of the nucleus
 Most of the mass of the atom is in the nucleus
 Most of the atom is empty space
Atomic number- the number of protons in a nucleus; symbol is Z
1. in a neutral atom, the number of protons is equal to the number
of electrons
Mass number is the number of protons and neutrons in a nucleus; symbol is A
1. mass number minus the atomic number equals the number of
neutrons in an atom
Isotopes: atoms that contain the same number of protons but a different number of
neutrons; they have different masses
1. represented by a nuclear symbol
2. Atomic mass/weight – is the average of all the naturally
occurring isotopes; expressed as amu; depends on the number
of isotopes and percent abundance
Examples:
1. How many protons, neutrons, and electrons are in 14C ?
2. How many protons, neutrons, and electrons are 11C ?
III.
The Periodic Table
A. You can identify elements by their atomic number, which can be read off
the periodic table.
B. History
1. Dimitri Mendeleev noticed a periodicity in the properties when
elements were arranged in increasing atomic mass
2. GJ Moseley reformed the hypothesis to state that periodic
properties are functions of increasing atomic number
C. Organization
1. Period- horizontal rows
2. Group – vertical columns also called families
a. Groups are numbered from 1-18 starting at left
1. group number of representative element or main
group elements are 1,2,13-18
2. transition elements are in the center of the periodic
table- groups 3 – 12; contain many elements with
more then 1 charge
b. Groups can also be numbered with a letter
1. A elements are called representative element or
main group
2. B elements are called transition- center elements;
all metallic solids
c. Groups with special names
1. Alkali metals: group 1 or 1A
a. properties- react vigorously with water to
produce hydrogen gas & alkaline solutions,
all metals, all solids, always combined,
easily lose 1 valence electron; form +1 ions
2. Alkaline earth metals: group 2 or 2A
a. properties- reactive but not as strong as 1A
all metals; all solids; form alkaline solutions,
lose 2 valence electrons, forms +2 ions
3. Halogens: group 17 or 7A
a. properties: very reactive nonmetals, react
with alkali metals to form salts, F,Cl, I, Br
are diatomic, gains 1 valence electron,
forms -1 ions
4. Noble gases (rare or inert): group 18 or 8A
a. properties- does not react with any other
substances, all gases, stable, does not form
ions
D. Metals, Nonmetals, and Metalloids
a. diagonal line or stair way starts to the left of boron
separates the metals from the nonmetals
b. metals are great conductors of heat &electricity, mostly
solids, high luster, malleable, ductile, and form alloys
c. nonmetals are poor conductors, dull & brittle, mostly
gases and liquids
d. metalloids lie to the immediate right or left of the stairway
with the exception of Al
1. share properties of metals and nonmetals
2. B, Si, Ge, As, Sb, Te, Po, and At are the 8
metalloids
VI Chemical Formulas
A.
Shorthand way of expressing compounds by indicating the type by symbol
and number of atoms by subscript
Molecular formula: shows the exact number of atoms of each element in the
smallest unit of a substance
 Molecules are 2 or more atoms of the same or different
nonmetal elements covalently bonded together; act as
discreet units
example: C6H6
 Diatomic Molecules contain only 2 atoms (Professor
BrINClHOF references the diatomic elements)
H2, N2, O2, Br2, F2, I2, Cl2, Br2, HCl, CO
 Covalent bonding occurs when valence electrons are shared
between 2 nonmetal atoms
 Forces between neighboring molecules are usually weak in
comparison to ionic bonding
Empirical formula : shows the simplest whole-number ratio of the atoms in a substance;
 Ions are charged particles formed when an atom or group of
atoms gain or lose an electron
 Cations are positively charged ions formed from metals that
lose electrons
example: Na+1
 Anions are negatively charged ions formed from nonmetals
that gain electrons example: Cl-1
 Ionic Bonding occurs when the oppositely charged ions are
attracted to each other: example NaCl
 Forces between neighboring formula units are usually
stronger than between molecules
V. The Nature of Ions
A. Number of protons and electrons in an ion
1. the number of protons remains unchanged
2. the number of electrons changes
a. For cations, the number of electron decreases by
the size of the charge, for example Li+1
b. For anions, the number of electrons increases by
the size of the charge, for example, F-1
Example
How many protons and electrons are in 27Al3+ ?
How many protons and electrons are in
78Se2-
?
B. Monatomic ions are single atoms that have lost or gained electrons
1. For main group elements, the charges of the ions formed by the
atoms can be predicted by applying this principle: Atoms that are close
to a noble gas( the nonmetals) form ions that contain the same number
of electrons as the neighboring noble gas atom.
2. Metals of group 1A –3A form positive ions with a charge equal to the
group number
3. Transition metals- no predictable pattern; have several different ion
charges
a. are in groups 3 –12
b. These cations typically have charges of +1, +2, or +3.
c. They ordinarily do not have a noble gas structure
4. Nonmetals form ions with a negative charge equal to (8 minus the group
number)
5. Some common charges of ions formed by the main group elements
a. group 1 : +1
b. group 2 : +2
c. group 16: -2
d. group 17: -1
e. Al in group 13: +3
f. N in group 15: -3
6. H can either gain or lose electrons; H+1 or H-1
7. Valence electrons- electrons in the outermost shell of an atom
C. Polyatomic Ions
1. 2 common polyatomic cations are NH4+1 and Hg2+2.
2. Polyatomic ions usually have 1 or more oxygen atoms. These are
collectively called oxoanions.
3. Memorize the name, formula and charge of the polyatomic. Most
polyatomic ions contain one or more oxygen atoms. Collectively, these
are called oxoanions.
a. Rules for naming oxoanions
1. The suffix -ate is used for the anion with the larger number of
oxygen atoms
2. The suffix -ite is used for the anion containing fewer oxygen
atoms.
3. The prefix hypo- is used when the oxoanion has more than 2
forms: it means the fewest oxygen atoms
4.The prefix per- is used when the oxoanion has more than 2
forms: it means it has the largest number of oxygen atoms.
VI. Naming and Writing Formulas for Ionic Compounds
A. The principle of electrical neutrality : the total positive charge of the
cations in the formula must equal the total negative charge of the
anions.
B. The formula for an ionic compound shows the simplest whole number
ratio between cation and anion. This is always an empirical formula.
C. Ions are arranged in a 3d network called a crystal lattice
D. Rules for Writing Formulas
1. Compounds have electrical neutrality. Na+1 and S-2 must be
written as Na2S since you need 2 positive charges to balance the
–2 charge on the S.
2. The positive ion is always written before the negative ion.
3. If 2 or more polyatomic ions are used in the formula, enclose the
polyatomic ion in parentheses and put the number of ions you need
outside the parentheses as a subscript.
4. Do not write the charge of the ion in the formula. For example
Examples: Write chemical formulas for the following:
aluminum oxide
calcium bromide
sodium carbonate
sodium sulfide is Na2S, not Na+2S-2
E. Rules for Naming Ionic Compounds
1. When a metal is involved, the name of the metal is used.
2. When the metal ion can have 2 different charges, the charge of the
ion is indicated by writing it in Roman numerals in parentheses
after the name of the metal. I.e. Cu+1 is copper (I)
3. When a nonmetal is involved, -ide is added as a suffix to the root
word of the nonmetal (usually the first syllable); i.e. bromine
becomes bromide
4. Polyatomic ions retain their name.
Examples: Name the following ionic compounds
BaCl2
K2O
Mg(OH)2
KNO3
VII. Naming and Writing Formulas for Covalent Compounds
A. Naming Binary Molecular compounds
1. Rule: The first nonmetals get its full name. The second nonmetal gets its
root word + - ide. Both nonmetals get a prefix denoting how many atoms are
used to make the compound. However, when only 1 atom is used, the prefix
mono is not attached in the first part of the binary compound.
For Example: CO2 is carbon dioxide
2. Prefixes: mono (1) di (2) tri (3) tetra (4) penta (5) hexa (6) hepta (7)
octa (8) nona (9) deca (10)
Examples Write the names for the following covalent compounds
HI
NF3
SO2
N2Cl4
NO2
B. Acids
1. Defined as a substance that yields H+1 ions when dissolved in water
2. When the gaseous binary molecules form acidic aqueous solutions, the -ide
suffix is change to -ic. and the prefix hydro- is used.
Example: Name the following acids
HCl
H2S
3. When a hydrogen ion combines with an oxoanion like PO4-3, an oxoacid is
formed. The prefixes of the oxoanion, per, and hypo remain unchanged. the
suffix -ate is changed to -ic and the suffix -ite is changed to -ous. i.e. HClO3 is
choric acid.
Example: Name the following acids
HNO3
H2CO3
H2SO4
Mixed Practice
1.Dinitrogen monoxide
2.Potassium sulfide
3.Copper (II) nitrate
4.Dichlorine heptoxide
5.Chromium (III) sulfate
6.Ferric sulfite
7.Calcium oxide
8.Barium carbonate
9.Iodine monochloride
1.BaI2
2.P4S3
3.Ca(OH)2
4.FeCO3
5.Na2Cr2O7
6.I2O5
7.Cu(ClO4)2
8.CS2
9.B2Cl4