Download Chapter 2 Atoms, Elements, Orbitals, and Electron Configurations

Chapter 2 Atoms, Elements, Orbitals, and
Electron Configurations
2.1 Atomic Theory
• The last chapter dealt with chemistry on a large-scale or
macroscopic level.
• Now we need to look at matter on a sub-microscopic, or atomic
level.
Questions about the fundamental nature of matter can be traced as far
back as the Greek philosophers.
• Aristotle believed that matter could be divided indefinitely.
• Democritus argued that there was a limit.
John Dalton proposed atomic theory in 1808.
− The work atom literally means "uncuttable."
− All matter is composed of atoms.
− The atoms of a given element differ from the atoms of all other
elements.
Chemical compounds consist of atoms combined in specific ratios.
Law of Definite Proportions – atoms combine in whole number
ratios to form compounds.
Water is always H2O, not H2.1O or HO2
Law of Multiple Proportions – The same kinds of atoms can
combine in more than one proportion. If you change the ratio, you
change the compound.
H2O is water but H2O2 is hydrogen peroxide
Chemical reactions change the way that atoms are combined in
compounds. The atoms themselves are unchanged.
2H2 + O2 2H2O
Chapter 2 Page 1 of 19
The Make-up of an Atom
Atoms are composed of tiny subatomic particles.
(There are many more particles (e.g., ferions, bosons, etc. that have been
identified, but these three still provide the core that is needed to
understand most chemical properties and transitions.)
The three fundamental subatomic particles are:
1. Protons – positively charged subatomic particles
2. Neutrons – electrically neutral subatomic particles
3. Electrons – negatively charged subatomic particles.
Characteristics of Particles
Particle
Mass (g)
Mass (amu)*
-24
Proton (p)
1.67262 x 10
1.007276 amu
-24
Neutron (n)
1.67493 x 10
1.008665 amu
-28
Electron (e ) 9.10933 x 10
0.000549 amu
-24
0.00091 x 10
*amu = atomic mass units to be explained later
Charge
+1
0
-1
Electrons are so much lighter that protons and neutrons that their mass is
usually ignored when computing the mass of an atom.
Atoms are extremely tiny.
Atomic diameters range from:
Hydrogen at 7.4 x 10-11 m
to
Cesium at 5.24 x 10-10 m
Atomic mass values range from:
Hydrogen at 1.67 x 10-24 g to
Uranium at 3.95 x 10-22 g
By comparison, a penny weighs 2.5 g
Chapter 2 Page 2 of 19
What is an amu?
The masses of subatomic particles are so small that they are hard to
work with.
• A relative scale was devised.
• Since carbon is one of the most abundant elements and it is the basis
for all living matter, the relative scale was based on carbon, C.
• Carbon-12 (aka C-12), which has 6 protons and 6 neutrons in its nucleus, is
the standard and is arbitrarily assigned the mass of 12amu.
• All other element's masses are defined relative to C-12.
1 amu = 1.6605 x 10-24g (1/12 the mass of one C-12 atom)
Problem: What is the mass, in grams, of a copper atom weighing
62.95 amu?
62.95 amu 1.6605 x10−24 g
x
= 1.045285 x10− 22 g ⇒ 1.045 x10− 22 g
amu
Problem: How many atoms are in 0.500g of sodium, Na, if each atom
of sodium weights 23 amu?
0.500 g
−
x
1 amu
1 Na atom
x
= 1.309 x 10 22 ⇒
−24
1.6605 x 10 g 23.0 amu
1.31 x10 22 Na atoms
−
Problem: What is the mass in grams of 2.4x1013 helium atoms, each
having a mass of 4.0 amu??
−
2.4 x1013 atoms 4.0 amu 1.6605 x 10 −24 g
x
x
= 1.6 x 10 −10 g
atom
amu
Chapter 2 Page 3 of 19
Arrangement of Subatomic Particles
The arrangement of the particles is not
random.
− Nucleus- a dense core within the atom
that contains the protons and
neutrons (and therefore almost all of
the mass).
− Surrounding space – is mostly empty
and contains the electrons.
− If an atom was the size of a large
domed stadium, the nucleus would
only occupy a volume the size of a
pea!!!
Structure of the Atom
• The structure of the atom is
determined by interplay of different forces.
• Opposite electrical charges attract each other.
• Like charges repel each other.
• The positively charged protons attract the negatively charged electrons.
- This keeps them in the area around the nucleus.
• Electron repel electrons – So they stay far away from each other.
- This accounts for the large space they occupy.
• Protons and neutrons in the nucleus are held together by the
nuclear strong force .
- This takes a lot of energy.
Chapter 2 Page 4 of 19
2.2 Elements and Atomic Number
• Atomic Number ( Z ) is the number of protons in atoms of a
given element.
– All atoms of a particular element have the same number of
protons in the nucleus.
– They will all have the same number of electrons , since the
atom is neutral.
– They may have different numbers of neutrons .
- If so they are isotopes of each other.
• Mass Number (A) is the sum of the protons plus neutrons in an atom.
2.3 Isotopes and Atomic Weight
Isotopes
Definition: Atoms (of the same element) with identical atomic numbers
but different mass numbers.
Representation:
Mass number →
AX
Atomic number → Z
← Element symbol
Z = atomic number = # of protons
A = Z(# of protons) + N (# of neutrons)
There are three isotopes of hydrogen:
−
1
1
H protium
2
1
H deuterium or D
Chapter 2 Page 5 of 19
3
1
H tritium or T
Most isotopes do not have specific names the way those of hydrogen do.
Isotopes are usually referred to by their mass number. For example:
235
92
U
is referred to as uranium-235 or U-235.
Problem: Which element has atomic number 17?
(ANS: Cl)
Problem: How many protons, neutron, and electrons are in the isotope
with Z=13 and A=27? What is the element?
a) # p = Z = 13
# e = Z = 13 (since it is neutral)
# n = A – Z = 27 – 13 = 14
b) Element # 13 = Al
Problem: How many neutrons are in U-238?
(ANS: n=238-92=146)
Problem: Complete each of the following isotope symbols:
11
5
?
=
11
5
B
56
?
Fe
=
56
26
Fe
Atomic Weight
Naturally occurring elements are mixtures of isotopes.
Atomic Mass values (aka atomic weights) listed for elements are the
weighted average mass of an element’s atoms.
• The individual masses of the naturally occurring isotopes and the
percentage ( as a fractional abundance) of each must be known.
Atomic weight = Σ[(isotope abundance) × (isotope mass)]
The Greek symbol Σ indicates sum of terms
Chapter 2 Page 6 of 19
Problem: Copper consists of only two isotopes with masses and
abundances of 62.95 (70.5%) and 64.96 (??%). Calculate the
atomic mass of Cu to the correct number of significant figures
from the data. [ANS: see periodic chart. (63.5)]
All percentages must add up to 100% so the other is 100-70.5=29.5%
Cu (62.95*0.705) + (64.96*0.295) =
44.380 + 19.163 = 63.543 ⇒ 63.5 amu/atom
2.4 The Periodic Table
History of The Periodic Table
• Only 10 elements have been known since the beginning of
recorded history. (Sb, C, Cu, Au, Fe, Pb, Hg, Ag, S, Sn)
• By the American Revolution, 24 elements had been identified.
• In the late 1700s and early 1800s, as more elements were
discovered, chemists began to look for similarities among
elements.
• Dmitri Mendeleev organized elements by mass, then into columns
based on chemical behavior.
• Holes in the table predicted the discovery of new elements.
Mendeleev’s
1871
Periodic Table
There are over 110 elements currently identified.
Chapter 2 Page 7 of 19
This is the periodic table in the front of your book. (Show on overhead.)
This is the most current one that you can always download. I like the format
better.
Most Current
Periodic Table from
webelements.com
This is the one we will be using in class. It has the format of webelements but
the same atomic weights as your textbook.
Chapter 2 Page 8 of 19
Using The Periodic Table
Each box in the periodic Table gives us
specific information about that element.
−
− Columns are called Groups of Families
(elements in the same group exhibit similar chemical behavior)
− Main Groups - The two groups on the far left ( 1 - 2 ) and the six
on the far right ( 13 - 18 ). (These are often designated with an A.)
−
Transition Metal Groups - Elements in groups ( 3 – 12 ).
−
Inner Transition Metals—
The 14 elements following lanthanum are the lanthanides ,
and the 14 elements following actinium are the actinides .
− Rows are called Periods (Tell why)
An example of a
periodic property
A graph of atomic radius versus atomic number shows a periodic riseand-fall pattern. The maxima occur for group 1A , elements, and the
minima occur for group
7A elements.
Problem: Identify the group 5 element in period 4 (ANS: V)
Chapter 2 Page 9 of 19
2.5 Characteristics of Different Groups
Periodic Table Groupings Handout
Group I (1A) – Alkali Metals
− Shiny, soft metals with low melting points.
− Not found in nature as the free element (very reactive)
− Form alkaline solutions in water
Group 2 (2A) – Alkaline Earth Metals
− Lusterous, silvery metals
Show sodium and potassium in
water movie. Reaction is:
2Na + 2H2O H2 + 2NaOH + heat
(requires Quicktime)
− Reactive (but less so than Group 1)
− Not found in nature as the free element
− Form alkaline solutions in water
Group 17 (7A) - Halogens
− Colorless and corrosive nonmetals
− Only found in nature combined with metals (e.g. NaCl)
− All except astatine (At) are diatomic
They exist as pairs as pure elements; i.e. F2, Cl2, Br2, I2.
(Note: 3 additional elements are diatomic: H2, N2, O2.
Group 18 (8A) – Noble Gases
− Colorless gases
− Very unreactive (previously known as inert gases)
(Groups 3-12) Transition Elements &
(no group number) Lanthanide/Actinides Series
− Transition elements are common
− Lanthanides and Actinides are not common
Chapter 2 Page 10 of 19
Zigzag Line
The periodic table can be grouped into 3 categories: Metals, nonmetals,
and metalloids (semimetals).
Metals
− Left side of the periodic table
− Most elements are metals
− They do not combine with each other; they combine with
nonmetals
− solid (except Hg)
Nonmetals
− Right side of periodic table
(Far right column referred to as the noble gases[very non-reactive];
next group referred to as the halogens[very reactive])
− Seven are diatomic (come in pairs in nature)
H2, N2, O2, F2, Cl2, Br2, I2
(Emphasize that they are only diatomic when they are the free
element, not when they are combined with other elements.)
− can be solids, liquids, or gasses
Metalloids
− Share some characteristics with both metals and non-metals.
− Are near the zigzag line from Boron to Astatine
(Metals to the left of the line, nonmetals to the right)
Textbooks differ; Semimetals are B, Si, Ge, As, Sb, Te,
(Sometimes astatine (At) and/or polonium (Po).
We will consider all elements touching the line (except Al) as a
semimetal.
Chapter 2 Page 11 of 19
2.6 Electronic Structure of Atoms
The properties of the elements are determined by the arrangement of
electrons in their atoms.
This arrangement is understood using the
quantum mechanical model developed by Erwin Schrödinger.
– (According to this theory), atoms are not completely free to move
around the atom.
• Each electron is restricted to specific probability regions
that are known as orbitals.
– The energy associated with each orbital is different
• The energy that electrons can have is quantized
(can only have discreet values)
e.g., stairs are quantized, a ramp is not.
– Electrons have both
properties.
particle
- like and
wave - like
– The behavior of electrons can be described using an equation
called a wave function.
• The wave function also provide each electron with an
“address” within an atom consisting of a shell, subshell,
and orbital.
Shells, Subshells, and Orbitals
Orbitals are grouped into shells and subshells.
Shells (aka Primary Quantum Number, n)
Electrons are grouped in layers or shells around the nucleus.
• The farther out from the nucleus,
- the larger the shell
- the more electrons it can hold, and
- the higher the energy level of the electrons.
Chapter 2 Page 12 of 19
Shell Numbering:
• Shells are numbered with integers starting with 1.
• A shell can have as many subshells as the shell value.
Subshells
Subshells “Numbering”:
Subshells are lettered in order s, p, d, f
Therefore:
- Shell n=1 has 1 subshell (only an s subshell)
- Shell n=2 has 2 subshells (s and p subshells)
- Shell n=3 has 3 subshells (s, p, and d subshells)
- Shell n=4 has 4 subshells (s, p, d and f subshells)
Leave room for 3 more rows!
Orbitals
Each subshell contains a grouping of orbitals.
s subshells can have 1 orbital
d subshells can have 5 orbitals
p subshells can have 3 orbitals
f subshells can have 7 orbitals
Each orbital can hold 2 electrons of opposite spin.
Chapter 2 Page 13 of 19
Orbital Summary
Problem: What is the maximum # of e- that can occupy the following
subshells?
The number does not affect the # of
3p? (6)
2s? (2)
2p? (6)
e- a subshell can hold!
Problem: What is the maximum # of electrons that can occupy the third
shell?
3s(2) + 3p(6) + 3d(10) = 18
Problem: How many electrons are present in an atom in which the 1s, 2s,
and 2p subsells are filled? Name the element.
1s(2) + 2s(2) + 2p(6) =10 e- ∴element = Ne
Orbital Shapes
Since orbitals are “probability regions”, they have shapes.
s orbitals are spheres.
p orbitals are dumbbells.
1 lobe
2 lobes
Chapter 2 Page 14 of 19
The shapes of d (4 lobes
per orbital) and f orbitals
(8 lobes per orbital) are
more complex.
The shapes of f subshells
are very complex.
Usually 8 lobes
Usually 4 lobes
2.7–2.8 Electronic Configurations and the Periodic Table
The exact arrangement of electrons in an atom’s shells and subshells is the
atom’s electron configuration. It can be predicted by applying 3 rules.
The 3 Principles for Order of Filling Energy Levels
• Rule 1: Always fill the lowest
levels first. (Aufbau Principle)
energy
- This is complicated by “crossover” of
energies above the 3p sublevel.
- The energy diagram can be used to predict
the order of filling, but there are easier
ways!
• Rule 2:Each orbital can hold a maximum of 2
electrons of opposite spin. (Pauli Exclusion
Principle)
• Rule 3: If you have orbitals of the same energy
( degenerate ), you fill 1 in each orbital before you double up. All
unpaired e- all have the same spin.
(Hund’s Rule)
Chapter 2 Page 15 of 19
The periodic table provides a method for remembering the order of
orbital filling. It can be divided into “blocks” elements according to the
last subshell filled.
Atomic Electron Configurations
A summary of the orbital location of each electron is written as, for
example:
Mg 1s2 2s2 2p6 3s2 etc.
Shell #
Subshell
Type
The number of
electrons in the
subshell
Chapter 2 Page 16 of 19
Problem: Determine the electron configuration and orbital box diagram for:
V
1s2 2s2 2p6 3s2 3p6 4s2 3d3
OR [Ar] 4s2 3d3
often re-written with the
OR [Ar] 3d3 4s2
This can be called
valence___electrons last (outer shell)
shortcut
notation or a
noble gas
notation.
Have class try: W
[Xe] 6s2 4f14 5d4 ⇒ [Xe] 4f14 5d4 6s2
An alternate trick that can be
used to determine the order or
sublevel filling is:
Chapter 2 Page 17 of 19
Problem: Determine the electron configuration for the following
elements. Also show the orbital diagram for each. Determine
the number of unpaired electrons.
Al 1s22s22p63s23p1
↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ .
1s 2s
2p
3s
V
ONE UNPAIRED ELECTRON
1s22s22p63s23p64s23d3
↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ ↑_ __ __
1s 2s
2p
3s
3p
4s
3d
THREE UNPAIRED ELECTRONS
Problem: Determine the noble gas electron configuration for the
following elements.
Se
1s22s22p63s23p64s23d104p4
[Ar] 4s23d104p4
Sn
1s22s22p63s23p64s23d104p65s24d105p2
[Kr] 5s24d105p2
Bi
1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p3
[Xe] 6s24f145d106p3
Chapter 2 Page 18 of 19
2.9 Electron Dot Symbols
Since valence electrons are so important in the behavior of atoms, it is
useful to be able to show them with the chemical symbol.
Electron-dot symbol—An atomic symbol with dots placed around it to
indicate the number of valence electrons.
These are cartoons and do not imply actual shapes.
Draw out row 1 on the board.
Problem: Write the electron dot symbol for element X in group 5 (5A).
Problem: Write the electron dot symbol for lead.
Problem: Write the electron dot symbol for strontium.
Chapter 2 Page 19 of 19