Download O 2 (g)

Lecture 4
Ch.3, 4.1-4.3
Suggested HW
Ch 3: 4, 28
Ch 4: 1, 4, 5, 12
Elemental Classifications
• Elements can be metals, nonmetals, or semiconductors (we
will discuss semiconductors later)
Physical Characteristics of Metal
•
•
•
•
•
•
Malleable
Ductile
Conductive of electricity
Conductive of heat
Have luster and shine
Very High Melting Points
Elemental Classifications
Physical Characteristics of Nonmetals
• Most nonmetals are gases
• Non conductive of heat and electricity
• Nonmetal solids are brittle, powdery
• Low melting points
O2 (g)
S(s)
He (g)
All elements LEFT of the black line are metals, except Hydrogen.
Chemical Groups and Diatomic Molecules
• Certain elements are unstable, and hence, do not
commonly exist as individual species, but as diatomic
molecules
• These include H, O, N, and all of the halogens (group 17)
H  H2 (Hydrogen gas)
O  O2 (Oxygen gas)
N  N2 (Nitrogen gas)
F, Cl, Br, I  F2, Cl2, Br2, I2
(Fluorine gas, chlorine gas, bromine gas and iodine gas)
Nomenclature: Ionic Compounds
• Molecules are formed by the atomic bonding. There are
two types of bonds: IONIC and COVALENT. Special rules
exist for naming molecules of each type.
• Ionic bonds form between metal ions and nonmetal ions
(further detail provided in chapter 6)
• To name an ionic compound, you do the following
1. Write the name of the metal
2. Follow it with the ionic name of the nonmetal
Example
• KF Potassium Fluoride
Nomenclature: Covalent Compounds
• Covalent bonds form between two
nonmetals (further detail provided in
chapter 7)
• To name an covalent compound:
1. Write the name of the first nonmetal.
For non-unity subscripts, use greek
prefixes (shown on right)
2. Follow that with the ionic name of the
second nonmetal. Again, include greek
prefixes. Only use mono- for oxygen
containing molecules.
Examples
CO carbon monoxide
CO2 carbon dioxide
N2S
P4Se10
dinitrogen sulfide
tetraphosphorus decaselenide
Nomenclature: Hydrogen
• Hydrogen is strange. It’s a nonmetal, but tends to react in a
manner similar to that the metals in group 1.
1. If hydrogen is listed first (not including halogens), use ionic
rules to name the molecule.
Ex. H2S = hydrogen sulfide
2. Hydrogen halides (HX where X is a halogen) are acids and are
named as such. We drop –gen and end the second nonmetal
with the suffix “–ic acid”
HCl = hydrochloric acid ; HF = hydrofluoric acid
3. If hydrogen is listed last, it is a hydride anion (H-)
Ex. MgH2 = magnesium hydride; TeH2 = tellerium dihydride
Group Work
• Name the following:
1.
2.
3.
4.
5.
SrO
IF3
HBr
CF4
NaH
Chemical Reactions
𝐍𝐚 𝐬 + 𝐂𝐥𝟐 𝐠 → 𝐍𝐚𝐂𝐥 𝐬 ?
• As you can see from the chemical
equation shown to the left, products
typically exhibit vastly different
characteristics the reactants
• Also recall our discussion on the law
of conservation of mass. Based on
this law, can you find a problem with
the equation written shown?
Balanced Reactions
• Mass can not be created or destroyed. This means that every
element contained in the reactants must be accounted for in
the product(s)
𝐍𝐚 𝐬 + 𝐂𝐥𝟐 𝐠 → 𝐍𝐚𝐂𝐥 (𝐬)
• There are two chlorine atoms on the reactant side, and only
one chlorine atom one the product side. To balance the
chlorine atoms, we add a coefficient of 2 to the NaCl(s)
𝐍𝐚 𝐬 + 𝐂𝐥𝟐 𝐠 → 𝟐 𝐍𝐚𝐂𝐥 (𝐬)
• We have balanced the chlorine atoms, but the sodium atoms
are now unbalanced. We add a coefficient of 2 to the Na (s).
The reaction is now balanced.
𝟐 𝐍𝐚 𝐬 + 𝐂𝐥𝟐 𝐠 → 𝟐 𝐍𝐚𝐂𝐥 (𝐬)
Coefficients And Subscripts
𝟐 𝐍𝐚 𝐬 + 𝐂𝐥𝟐 𝐠 → 𝟐 𝐍𝐚𝐂𝐥 (𝐬)
• The balanced equation above says that two Na atoms react with
one chlorine gas molecule to produce two molecules of NaCl
• The coefficient of 2 means that there are two separate Na atoms
• The subscript of 2 indicates two Cl atoms bonded together in a
single molecule
Na
Na
Cl
Cl
NaCl
NaCl
• Do not confuse coefficients and subscripts. Do not alter
subscripts when balancing.
Tips For Balancing Reactions
• Before carrying out any calculations, it is imperative that
you first confirm that a given chemical equation is
balanced.
• The rules for balancing a chemical equation are provided
below.
1. First, balance those elements that appear only once
on each side of the equation
2. Balance the other elements as needed. Pay attention
to subscripts.
3. Include phases
Tips For Balancing Reactions
• Let’s balance the equation below using the rules from the
previous slide.
C3H8 (s) + O2 (g)
CO2 (g) + H2O (L)
• We’ll balance C first.
C3H8 (s) + O2 (g)
3 CO2 (g) + H2O (L)
• Now balance H.
C3H8 (s) + O2 (g)
3 CO2 (g) + 4 H2O (L)
• Now balance O.
C3H8 (s) + 5 O2 (g)
3 CO2 (g) + 4 H2O (L)
Group Work
Balance the following:
1. Sulfur (s) + Oxygen gas (g)  Sulfur trioxide (g)
2. Nitrogen gas (g) + Hydrogen gas (g)  NH3 (ammonia gas)
3. C4H10 (L) + O2 (g)  CO2 (g) + H2O (g)
A
B
C
n=4
n=3
n=2
n=1
Chemical Groups And Periodicity
• As scientists first began to discover and classify the elements,
patterns and similarities were observed in chemical behaviors
of certain groups of elements.
• Consider the three metals Li, Na, and K
– All 3 metals are soft
– All 3 metals are less dense than water
– All 3 metals have similar appearance and low melting points
– The most interesting feature is that all 3 metals react with
the same elements in a nearly identical manner
• As you see in the periodic table, these elements are all listed
in the same group, or vertical column.
Chemical Groups And Periodicity
• Dmitri Mendeleev created the periodic table in in 1869 by
arranging the elements from left to right in order of increasing
atomic number, and vertically according to their behavior
(groups)
• In doing so, he observed repetitive patterns in chemical
behavior across periods (horizontal rows)
• This periodicity is described in the next slide.
Chemical Groups And Periodicity
Decreasing metallic
character
3
Li
Highly reactive,
highly conductive
metal
4
Be
6
C
Less reactive,
less conductive
metal
Nonconductive,
nonmetallic
solid
9
F
Highly reactive,
diatomic,
nonmetallic gas
Decreasing metallic
character
11
Na
Highly reactive,
highly conductive
metal
19
K
Highly reactive,
highly conductive
metal
14
Si
12
Mg
Less reactive,
less conductive
metal
20
Be
Less reactive,
less conductive
metal
Slightly
conductive
semi-metal
Decreasing metallic
character
22
Ge
Slightly
conductive
semi-metal
17
Cl
Highly reactive,
diatomic,
nonmetallic gas
25
Br
Highly reactive,
diatomic,
nonmetallic liq.
10
Ne
Totally
unreactive
gas
18
Ar
Totally
unreactive
gas
26
Kr
Totally
unreactive
gas
WHY?
• We must now answer many questions about chemical
reactivity.
– Why is it that some atoms join together and form molecules,
while others can’t?
– Why is there such wide variation in the reactivity and
physical properties of elements?
– Why is there periodic repetition (periodicity) of the
chemical/physical properties of elements as we move
across the periodic table?
Explanation Of Elemental Groups
• As previously discussed, Mendeleev noticed that chemical
behavior was repeated periodically when elements were
sorted by increasing atomic number
• The existence of periodicity proves a very important point:
The number of protons in the nucleus has no effect on
chemical behavior. If it were so, chemical behaviors would
never repeat given that no two elements have the same
atomic number.
The chemical behavior of an element must be dictated by
the configuration of electrons around the nucleus.
Ionization Energy
•
A direct indication of the arrangement of electrons about a
nucleus is given by the ionization energies of the atom
•
Ionization energy (IE) is the minimum energy needed to remove an
electron (form a cation) completely from a gaseous atom
– Ionizations are successive.
– As you remove one electron, it becomes increasingly difficult
to remove the next because of the increasing attraction
between the remaining electrons and the protons in the
nucleus
𝑀 → 𝑀+ + 𝑒 −
1st Ionization Energy
𝑀+ → 𝑀2+ + 𝑒 −
2nd Ionization Energy
IE1 < IE2 < IE3 …….IEn
Ionization Energy
• By measuring the energy required to remove electrons
from an element, you can gain an idea of how “willing” an
atom is to lose an electron, and relate this to its reactivity
• In the next slide, you will see data from an experiment in
which the 1st ionization energies of elements are plotted
against atomic number.
1st Ionization Energy Shows A Periodic Trend For T
very difficult to ionize
very easy to ionize
Ionization Energy
• It is relatively easy to
remove electrons from
group 1 metals.
– It becomes increasingly
difficult as you move
right across the
periodic table, and up a
group.
• It takes a very large amount
of energy to ionize a noble
gas.
• Like chemical properties,
ionization energies are also
periodic.
The lower the ionization
energy of an element, the
more METALLIC and
REACTIVE it is.
Electron Arrangement (Electronic Structure)
• The closer an electron is to the nucleus, the harder it
would be to pull the electron away.
– By carrying out multiple ionizations, we can gain
insight into the arrangement of electrons around the
nucleus of the element.
Example
• Using the table of ionization energies in the previous
slide, calculate the energy required to ionize Be to Be3+
• In order to go from Be to Be3+, you must LOSE 3
electrons. This will require 3 ionization steps (see
pg 107 in book).
𝐵𝑒 𝑔 → 𝐵𝑒 + 𝑔 + 𝑒 −
𝐼𝐸1 = 1.49 𝑎𝐽
𝐵𝑒 + 𝑔 → 𝐵𝑒 2+ 𝑔 + 𝑒 −
𝐼𝐸2 = 2.92 𝑎𝐽
𝐵𝑒 2+ 𝑔 → 𝐵𝑒 3+ 𝑔 + 𝑒 −
𝐼𝐸3 = 24.7 𝑎𝐽
29.1 aJ
Remember, energy is always in Joules (J). atto (a) = 10-18
Successive Ionizations
• Let’s take a look at the electron configurations of Lithium
(atomic # = 3) and Beryllium (atomic # = 4)
Li
3 electrons
Single
electron
that is easily
removed
Be
4 electrons
Pair of tightly
bound
electrons
Pair of
electrons that
are more easily
removed
Successive Ionizations
Ne
10 electrons
Na
11 electrons
Eight electrons of
similar attraction to
the nucleus
Same two tightly
bound electrons
11th electron
enters different
“shell”
Electrons Reside In “Shells” Of Different Distances
From The Nucleus
• From these plots, Niels Bohr derived the Bohr model of the
atom. In it, electrons reside in shells that orbit at different
distances from the nucleus.
• Each shell has a finite number of electrons that it can hold
• The two electrons closest to the nucleus are the hardest to
remove.
Na
• Each shell holds 2n2
electrons, where the
n=1 shell is the closest
to the nucleus.
Same Outer Electron Configuration Along A Group
Leads to Similarities in Reactivity
Na
Li
K
All group 1 metals have 1 lone electron
in the outermost occupied shell (valence
shell). Higher energy shells exist, but
are empty!
Chemical properties of an
element are determined by the
outer electron configuration.
Periodicity is Due To Repeating Valence Electron
Configurations
Li
Be
B
C
N
O
F
Ne
Na
Mg
Al
Si
P
S
Cl
Ar
Noble Gas Configurations
• The inner-most electrons of an element comprise what is
known as a noble gas core.
– At the close of each shell, you have a noble gas
configuration. Noble gases are chemically inactive
because they have completely filled shells.
• Lithium, for example, has a two electron core, which we call a
Helium core, and one outer, or valence electron. Sodium has a
10-electron, Neon core, and one valence electron; and so on.
• The electron configuration of an element can be represented
with a Lewis dot formula
Full Lewis dot
configuration
Valence Lewis dot
configuration
• We use these
representations to
describe the electron
configurations of an
element.