Download (+) = # of electrons

The Atom
Chapter 2
Chem2A
1
2
3
Elements
• An element is a pure substance that cannot be
broken down into simpler substances by a
chemical reaction.
• Consists of a dense core, surrounded by a
negatively charged cloud
• Contains three types of subatomic particles
– Proton
– Neutron
– Electron
4
The Proton
• Charge = +1.602x10-19 C
– +1
• Located in the nucleus
• Mass = 1.672 x 10-24 g
– ~1 atomic mass unit (amu)
– 1 amu = 1.66 x 10-24 g
• 1/12 the weight of a 12C atom
5
The Neutron
• No charge (O C)
• Located in the nucleus
• Mass = 1.675 x 10-24 g
– ~1 amu
6
The Electron
• Charge = -1.602 x 10-19 C
– -1
• Located outside the nucleus
in an e- “cloud”
• Mass = 9.109 x 10-28 g
– ~0 amu
7
Plum Pudding Model
• J.J. Thomson
• 1904
8
Gold Foil Experiment
• Ernest Rutherford
• 1909
9
Thomson vs. Rutherford
• Rutherford used the angle of deflection to calculate size of
the nucleus
• Nucleus approximated at 10,000 times smaller than the
radius of the atom
• Atom is actually mostly empty space
10
Atoms
11
12
13
Periods
14
Groups
15
Periodic Table of Elements
16
Classification of Periodic Table
17
Elements and Symbols
•
•
•
•
•
C = Carbon
N = Nitrogen
O = Oxygen
Cl = Chlorine
Ba = Barium
•
•
•
•
•
U=?
Cf = ?
Bk = ?
Md = ?
Es = ?
•
•
•
•
K=
Pb =
W=
Sb =
•
•
•
•
•
Co =
Cu =
Cr =
Mg =
Mn =
Metals
• Shiny
•Conduct Electricity
•Ductile
– Can be drawn into a thin wire
•Malleable (Shapeable)
•High M.P. and B.P.
•Solids @ RT
– Except Hg
20
Non-Metals
•
•
•
•
•
Don’t conduct well
Not usually ductile
Brittle
Low M.P. and B.P.
Many are gasses at
RT
21
Metalloids
• Chemical characteristics
in between metals and
non-metals
• Includes all atoms falling
on the dividing line
between metals and non
metals
22
Classification of Periodic Table
23
Alkali Metals
•
•
•
•
•
Group 1 (1A)
Li, Na, K etc.
Soft, shiny metals
Conduct heat and electricity
React violently with H2O
– Form H2(g) and alkaline
(basic) solutions
Akali(ne) Earth Metals
•
•
•
•
Group 2 (2A)
Be, Mg, Ca etc.
Not as reactive as Alkali Metals, but still quite reactive
Tend to make basic solutions when placed in water
Transition
Metals
• Groups 3-12
• Tend to have high
densities and B.P.
• All are metals
• Often used for
electrical conduction
• Often have vivid
colors when in
solution
– Used for pigments
Colors of Transition Metal
Compounds
Iron
Cobalt
Nickel
Copper
Zinc
Lanthanides
• Elements 57-71
– Lanthanum (La) to
Lutetium (Lu)
• Commonly used in lasers
• Can deflect UV and infrared
rays
Actinides/Actinoids
• Elements 89-103
– Actinium (Ac) to
Lawrencium (Lr)
• Only Actinium, Thorium (Th),
and Uranium (U) occur
naturally
– Others created by neutron
bombardment
• Radioactive
Groups 13(3A) – 16(6A)
• No common name
• Boundary between metals and non-metals
occurs here
• Contain elements abundant in earth’s crust,
atmosphere, and living things
• Contains the metalloids
Halogens
• Group 17 (7A)
• Love to form salts
with metals
– NaCl, KBr, CaCl2
• Like to form
diatomic molecules
– F2, Cl2, Br2
Noble Gases
• Group 18 (8A)
• Very unreactive
• Don’t like to bond to
other molecules
• Generally not
abundant
Diatomic Molecules
Dinitrogen (N2)
• Molecules consisting of only two atoms of either
the same or different elements
– O2
– CO
• Homonuclear Diatomic Molecule: a molecule
made of two atoms of the same element
– H2
• Heteronuclear Diatomic Molecule: a molecule
made of two atoms that are different elements
– NO
Allotropism
• The existence of multiple pure forms of an
element, in the same phase (solid, liquid, or gas),
that differ in structure
• Different forms are called allotropes
• Can exhibit varied physical properties and
chemical behaviors
• Don’t confuse allotropes with isotopes!!!
O2
O3
2.2 Structure of the Atom
From the periodic table:
Atomic Symbol
Atomic Number
= # of protons
Atomic Mass
• Any given element is neutral
• # of protons (+) = # of electrons (-)
• # of neutrons = Atomic Mass (rounded) - # Protons
– 7 – 3 = 4 neutrons
36
Problems
H
C
Fe
Pd
W
Atomic Weight
# Protons
# Electrons
# Neutrons
37
Calculator Boot Camp
38
2.3 Isotopes
A. Isotopes, Atomic Number, and Mass Number
Isotopes are atoms of the same element that have
a different number of neutrons.
the number of protons (Z)
+
the number of neutrons
Mass number (A) =
Mass number (A)
35
Atomic number (Z)
17
# of protons
Cl
37
17
Cl
Mass number (A)
Atomic number (Z)
=
17
# of protons
=
17
# of electrons =
17
# of electrons =
17
# of neutrons =
35 – 17 = 18
# of neutrons =
37 – 17 = 20
39
Identify the atomic mass, number of protons, and number
of neutrons for the following elements.
2.3 Isotopes
B. Atomic Weight
The atomic weight is the weighted average of the
masses of the naturally occurring isotopes of a
particular element reported in atomic mass units.
From the periodic table:
6
C
12.01
atomic number
element symbol
atomic weight (amu)
41
2.3 Isotopes
B. Atomic Weight
HOW TO Determine the Atomic Weight of an Element
Example
What is the average atomic weight of
chlorine?
Step [1]
List each isotope, its mass in atomic
mass units, and its abundance in nature.
Isotope
Mass (amu)
Isotopic Abundance
Cl-35
34.97
75.78% = 0.7578
Cl-37
36.97
24.22% = 0.2422
42
2.3 Isotopes
B. Atomic Weight
HOW TO Determine the Atomic Weight of an Element
Step [2]
Multiply the isotopic abundance by the mass
of each isotope, and add up the products.
The sum is the atomic weight of the element.
43
1. Using the percent abundances below, calculate the average atomic mass for
Carbon
a)
C-12 = 98.890%
b)
C-13 = 1.110%
2. Antimony (Sb) has two stable isotopes, 121Sb and 123Sb with masses of
120.9038 u and 122.9042 u, respectively. Calculate the percent abundances of
these two isotopes
45
Atomic Orbitals and
Electron Configurations
Electronic Structure
•Electrons are confined to discrete regions
called shells
•The shells are numbered n = 1, 2, 3, etc.
•Electrons in lower numbered shells are closer to
the nucleus and are lower in energy.
•Electrons in higher numbered shells are further
from the nucleus and are higher in energy.
47
Electronic Structure
• Shells with larger numbers (n) are farther from the
nucleus and can hold more electrons.
• # of electrons in each shell is calculated using
2(n2), where n = shell number.
Shell (n)
increasing
energy
Number of Electrons
in a Shell
4
32
3
18
2
8
1
2
increasing
number of
electrons
48
Electronic Structure
• Shells are divided into subshells, identified by the
letters s, p, d, and f.
• The subshells consist of orbitals.
• An orbital is a region of space where the
probability of finding an electron is high.
• Each orbital can hold two electrons.
increasing
energy
Subshell
s
p
Number of Orbitals
1
3
d
5
f
7
49
S Orbitals
• Spherical shape
• Lowest energy of the orbitals
• Contains 1 orbital: s
– 1 orbital can hold 2 electrons MAX
– = 2 electrons max
50
P Orbitals
• The p orbital has a dumbbell shape
• Higher energy than S Orbitals
• 3 separate orbitals px, py, pz
– Each orbitals can hold 2 electrons MAX
– = 6 electrons max
51
D Orbitals
• Double dumb bell shape
or single dumb bell with
a donut
• Higher energy than s
and p
• 5 orbitals
• dxy, dxz, dyz, dx2-y2, dz2
• = 10 electrons max
52
F Orbitals
•
•
•
•
Complex Shapes
Highest Energy Orbital
7 Orbitals
=14 electrons max
53
54
Electronic Structure
55
Electron Configuration
•
The electron configuration shows how the electrons
are arranged in an atom’s orbitals.
• The ground state is the lowest energy arrangement.
Rules to Determine the Ground State Electronic
Configuration of an Atom
Rule [1] •Electrons are placed in the lowest energy
orbital beginning with the 1s orbital.
•Orbitals are then filled in order of
increasing energy.
56
Electron Configuration
Rules to Determine the Ground State Electronic
Configuration of an Atom
57
Electron Configuration
Rules to Determine the Ground State Electronic
Configuration of an Atom
Rule [2] Each orbital holds a maximum of 2 electrons.
Rule [3] When orbitals are equal in energy:
•1 electron is added to each orbital until all
of the orbitals are half-filled.
•Then, the orbitals can be completely filled.
58
Electron Configuration
Orbital Diagrams
An orbital diagram uses a box to represent each
orbital and arrows to represent electrons.
an orbital
a single,
unpaired
electron
an electron
pair
Two electrons must have paired spins (opposite
directions) to fit into the same orbital.
59
Electron Configuration
A. First-Row Elements (Period 1)
Element
Orbital
Notation
Electron
Configuration
H (Z = 1)
1 electron
He (Z = 2)
2 electrons
60
Electron Configuration
B. Second-Row Elements (Period 2)
Element
Orbital
Notation
Electron
Configuration
Li (Z = 3)
3 electrons
C (Z = 6)
6 electrons
Ne (Z = 10)
10 electrons
61
Problems
• Determine the electron configuration and
orbital diagrams for the following atoms
1. N
2. O
3. Na
4. Mg
5. Fe
6. W
62
Electron Configuration
• The electron configuration can be shortened by
using Noble Gas Notation.
• Write the Symbol of the previous Noble Gas, then add
the electronic configuration of the additional
electrons.
element:
C
previous
noble gas:
63
Electron Configuration
C. Other Elements
Element
Ca
20 electrons
64
Electron Configurations
and the Periodic Table
65
2.7 Valence Electrons
• The chemical properties of an element depend
on the number of electrons in the valence shell.
• The valence shell is the outermost shell (the
highest value of n).
• The electrons in the valence shell are called
valence electrons.
Be
Cl
1s22s2
1s22s22p63s23p5
valence shell: n = 2
valence shell: n = 3
# of
valence electrons = 2
# of
valence electrons = 7 66
2.7 Valence Electrons
A. Relating Valence Electrons to Group Number
• Elements in the same group have similar
electron configurations.
• Elements in the same group have the same
number of valence electrons.
• The group number, 1A–8A, equals the
number of valence electrons for the main
group elements.
• The exception is He, which has only 2
valence electrons.
• The chemical properties of a group are
therefore very similar.
67
2.7 Valence Electrons
A. Relating Valence Electrons to Group Number
Group number:
1A
2A
3A
Period 1:
H
1s1
4A
5A
6A
7A
8A
He
1s2
Period 2:
Li
2s1
Be
B
C
N
O
F
Ne
2s2 2s22p1 2s22p2 2s22p3 2s22p4 2s22p5 2s22p6
Period 3:
Na
Mg
Al
Si
P
S
Cl
Ar
3s1 3s2 3s23p1 3s23p2 3s23p3 3s23p4 3s23p5 3s23p6
Valence Electrons
B. Electron-Dot Symbols
• Dots representing valence electrons are placed
on the four sides of an element symbol.
• Each dot represents one valence electron.
• For 1 to 4 valence electrons, single dots are
used. With more than 4 valence electrons, the
dots are paired.
Element:
H
C
O
Cl
# of Valence electrons:
1
4
6
7
Electron-dot symbol:
H
C
O
Cl
69
Periodic Trends
A. Atomic Size
Increases
•The size of atoms
increases down a
column
Decreases
•The size of atoms decreases across a row
70
Chapter 2 Review
• The periodic table will be folded up for quiz and
exam
• Know the names and elemental symbols of s
block, p block, and select d block atoms
– W, Cu, Ag, Fe, Hg, Ti
71
Chapter 2 Review
• Know how to classify atoms as metals,
non-metals, and metalloids.
• Know the period classifications
– Alkali, Alkaline, Transition Metals, main block,
halogens, noble gas
72
Classification of Periodic Table
13
Al
84
Po
73
Chapter 2 Review
• Location, and charge of the 3 particles that
make up an atom
– Proton (nucleus), Neutron (nucleus), Electron
(diffuse cloud surrounding nucleus)
• Determining how many protons, neutrons
and electrons an atom has
74
Chapter 2 Review
• Orbitals
– s, p, d, f
– Know everything about s and p
– Know how many electrons all can have
– Know how they relate in energy
75
Chapter 2 Review
• What are isotopes
• How do isotopes relate in terms of weight
• How to calculate the average atomic
weight
76
Chapter 2 Review
• Know how to fill in (yes ill give you boxes)
molecular orbitals
• Know how to write noble gas abbreviations
for any atom
77
Chapter 2 Review
• Know what valence electrons are and how
to ID them in an electron configuration
• Know the relationship between atoms in
the same period
• Know how to write electron dot symbols
78