Download atomic theory - unit a

CHAPTER 2. ATOMS
2.2 COMPOSITION OF MATTER
1. Pure substance: Consist of 1 Substance with fixed composition and distinct
properties; cannot be separated by physical means
A. Element: Made up of unique atoms & cannot be chemically separated
into simpler substances. E.g. H, O, C
B. Compound: Consist of 2 or more different elements & can
be chemically broken down into simpler substances. E.g. H2O, NaCl
• Compounds have different properties than elements contained in compound.
E.g. H2O vs. H2 & O2
2. Mixture: Physical combination of 2 or more substances; variable composition;
components can be separated by physical means
E.g. sand, rocks, egg, salt water, air, vodka tonic
ELEMENTS & COMPOUNDS
• 112 known elements (83-112 are radioactive)
• ~88 naturally occurring elements
• Each element has a 1 - 3 letter symbol derived from its name; 1st letter & only
1st letter is always capitalized.
*
Memorize names & symbols of the selected elements on handout
Chemical formula of an element or compound tells us type & number of atoms
present. E.g. Ne, H2 , H2O, Ca(OH)2
• subscript indicates how many atoms are present; 1 is omitted
2.3 ATOM: smallest particle of an element that retains the chemical identity of the
element (John Dalton).
• Atomic size: 0.1 - 0.5 nm
• 1981 - STM (scanning tunneling microscope) used to "see" atoms
Dalton's Atomic Theory - 1808
1) Each element is composed of small particles called atoms.
2) All atoms of an element are identical to each other.
3) Atoms cannot be created or destroyed.
4) Atoms of one element are different than atoms of another element.
5) Compounds are formed when atoms of different elements combine.
molecule: 2 or more atoms bonded together; the atoms may be identical or
different from one another.
•
Many elements exist as diatomic molecules: O2, N2, H2, F2, Cl2, Br2, I2
Law of Constant Composition: A compound always consists of the same
combination of elements (% composition is fixed).
E.g. 100 g of H2O always contains 11.1 g H, 88.9 g O
Law of Conservation of mass: during a chemical reaction, mass is conserved.
Total mass of the reactants = Total mass of the products
2.4 SUBATOMIC PARTICLES
•
The 3 main subatomic particles are the proton, electron and neutron:
Particles
electrons
protons
neutons
charge
-1
+1
0
mass (g)
mass (amu)
-28
9.107 x 10
g = 1/1835 amu
-24
1.672 x 10
g = 1 amu
1.675 x 10-24 g = 1 amu
location
outside nucleus
nucleus
nucleus
Nuclear Model (Rutherford):
1) Each atom has a very small nucleus. Protons & neutrons are located in the small
nucleus; it contains most of the mass.
2) Electrons are located in region outside of the nucleus. Most of the atom is empty
space.
3) Atoms possess neutral overall charge - equal number of protons & electrons.
ATOMIC NUMBER, MASS NUMBER, ISOTOPES
•
•
•
Every atom of an element has the same # of protons
The # of protons defines an element
The number of protons is called the atomic number, Z. (top # on per. table)
Mass number = A = total number protons & neutrons in nucleus
A = number p's + number n's
n=A-Z
Isotopes are atoms of an element that have a different number of neutrons.
Isotopes of an element have the same atomic number, but a different mass
number.
•
•
•
For neutral atoms: # of protons = # of electrons
nuclear symbol(isotope symbol) for element X:
Name of isotope: name-A
Ex.
35
17
Cl
Ex.
66
30
Zn
•
Z = p = 17
e- = 17
Z = 30
A = 66
A
Z
X
A = p + n; Z = p
A = 35
n = A - Z = 35 - 17 = 18
name: chlorine-35
p = 30
e- = 30
n = 36
name: zinc-66
Note that Z is sometimes omitted from the nuclear symbol:
13
13
C can be used for carbon-13
6 C or
IONS - atoms that have lost or gained electrons. Charge is written after symbol.
Ions are discussed in more detail in Chapter 3.
E.g. Mg+2 ion has 12 protons and 10 electrons (positively charged ion has more p)
F- ion has 9 protons and 10 electrons (negatively charged ion has more e-)
ATOMIC MASS
Masses of atoms are so small that we define the atomic mass unit (amu) to scale up
the numbers
•
The carbon-12 isotope was assigned a mass of exactly 12 amu, masses of other
elements are scaled relative to C-12.
•
1 amu = 1.66054 x 10-24 g, so p & n have masses of approximately 1 amu.
Atomic mass of an element is the average weighted mass of all isotopes of an
element. (Atomic mass is found at bottom of box on periodic table)
2.5 THE PERIODIC TABLE
•
1st table 1869 produced by Mendeleev & Meyer
Features of modern periodic table:
1. Elements arranged in order of increasing Z (atomic #), not mass.
2. Horizontal Rows in periodic table are called periods. 7 periods exist
3. Vertical Columns on table correspond to groups or families; elements in a
group have similar properties.
representative elements: A Group
transition elements: B Group
4. Metals are located to the left of the stairstep line. Metals are shiny, malleable,
ductile, good conductors, solids (except Hg)
5. Nonmetals are located to the right of the stairstep. Nonmetals are dull,
insulators, can be gas, liquid or solid
6. Elements located at the diagonal line are intermediate in character. These are
semiconductor or metalloid elements. Si, Ge, As, Sb, Te, Po, At
•
Individual boxes on periodic table have atomic # at top, chemical symbol in
middle, atomic mass at bottom
E.g. As
Atomic # = z = 33; As = arsenic;
mass = 74.92 amu
2.6 Electronic Stucture
Bohr model ~1913
1) charged e's travel rapidly in orbits around the tiny + charged nucleus
2) Electrons are contained in specific energy levels called principle energy levels or
shells. These energy levels are quantized which means only certain energies are
allowed. There is not a continuum of energy.
Analogy: stairway (quantized)
ramp (continuum)
Energy levels are designated by the quantum number n. n = 1,2,3….
n = 1 is ground state level - this is level closest to nucleus.
3) Electrons can jump from one level to another by absorbing or emitting energy.
QUANTUM MODEL - SCHRODINGER
1) n = principal quantum number, where n is energy shell.
Values for n = 1,2,3,4...
(n = 1 closest shell to nucleus)
Generally, energy increases with increasing n.
2) Principle energy levels can be subdivided into subshells.
• Electrons within a subshell have identical energy.
• There are 4 known subshells: s, p, d, and f. These correspond to blocks of
elements on periodic table.
3) Each subshell contains a specific number of orbitals
Orbitals are regions in space where an electron is most likely to be found.
Sublevel
s
p
d
f
Block on Chart
1A, 2A
3A-8A
transition metals
inner transition
Max # of e2
6
10
14
# of orbitals
1
3
5
7
Electron Configurations and Orbital Diagrams
Important Rules
1) Pauli exclusion principle - no more than 2 e- can occupy the same orbital
2) Aufbau principle: Orbitals are filled in order of increasing energy. At ground
state, electrons fill lowest energy orbitals available
3) Hund's rule: When filling degenerate orbitals (same subshell), 1 e- is placed in
each orbital before pairing up.
An electron configuration describes how many e- an atom has in each of its
subshells.
Suggestions for writing electron configurations
1) Locate element on periodic table.
2) List all subshells in order of increasing energy.
3) Use superscript #'s to indicate # of e-'s in each subshell.
4) Double check. Sum of superscript #'s = Z
Be able to write electron configurations for 1st 20 elements
E.g. Write electron configurations & orbital diagrams for the following
H (Z = 1)
1s1
He (Z = 2) 1s2
Li (Z = 3)
1s22s1
C (Z = 6)
1s22s22p2
F (Z = 9)
1s22s22p5
Mg (Z = 12) 1s22s22p63s2
S (Z = 16) 1s22s22p63s23p4
Noble gas core - can use noble gas from previous row as short cut.
Na (Z = 11)
1s22s22p63s1
[Ne]3s1
!
[Ne]
Cl (Z = 17)
[Ne]3s23p5
2.7 PERIODIC PROPERTIES
VALENCE ELECTRONS
•
•
The outer shell s and p electrons are called the valence electrons. These
electrons dictate the properties of an element.
All elements in a family have identical valence shell e-configurations - same # of
valence electrons.
Group
e- conf.
valence e-
1A
2A
3A
ns1
ns2
ns2np1
1
2
3
4A
ns2np2
4
5A
ns2np3
5
6A
ns2np4
6
7A
ns2np5
7
8A
ns2np6
8
Ionization Energy - energy required to remove an electron from an atom.
•
IE increases → periodic table
- Harder to remove e-'s from Noble gases & nonmetals (don't want to lose e-'s)
•
IE decreases ↓ periodic table
- Easier to remove e- from higher energy shell (further from nucleus)