Download Unit 4 Notes - Vinita Hornets

CHEMISTRY:
ATOMS, MOLECULES &
REACTIONS
HAGEBUSCH
NEO A&M – PHYS 1014
Atomic Theory
Atom – smallest part of an element that still
retains the properties of that element
Atomic Theory
Dalton’s Atomic Theory:
1. All elements are composed of atoms
2. All atoms of an element are identical and differ
from atoms of other elements
3. Atoms of different elements can chemically
combine in simple whole number ratios
4. Chemical reactions occur when atoms are
separated, joined, or rearranged.
 Atoms are never turned into a different atom
during a reaction
Atomic Structure
Electron (e-) – subatomic particle with a negative
charge
• Discovered by English physicist JJ Thompson,
1897
• Smallest of all subatomic particles
• Electrons are located in empty space around
the nucleus called an electron cloud.
The Electron

Streams of negatively charged particles were
found to emanate from cathode tubes.
The Atom, circa 1900


The prevailing theory
was that of the “plum
pudding” model, put
forward by Thompson.
It featured a positive
sphere of matter with
negative electrons
imbedded in it.
Rutherford’s Gold Foil Experiment
Rutherford’s Gold Foil Experiment
Atomic Structure
Atomic Nuclei - central part of an atom that
contains about 99.9% of atomic mass
1. Proton (p+) – subatomic particle with a positive
charge
 Discovered
by Rutherford, 1919
2. Neutron (no) – subatomic particle with no charge
 Discovered
by James Chadwick, 1932
Atomic Structure
The Bohr Model of the Atom
• Neils Bohr, a Danish physicist, proposed that
electrons occupy specific orbits around the
nucleus called energy levels
• Energy level – region around the nucleus of an
atom where an electron is likely to be moving
Atomic Structure
The Bohr Model of the Atom
 He proposed that the electrons have a fixed
energy and that keeps them away from the
nucleus
 The closer the energy level is to the nucleus,
the lower the energy of the electrons in it
 Bohr discovered this by studying the emission
spectrum emitted by atoms after they were
charged with high voltage electricity
Subatomic Particles



Protons and electrons are the only particles that
have a charge.
Protons and neutrons have essentially the same
mass.
The mass of an electron is so small we ignore it.
Distinguishing Between Atoms
Atomic Number – the number of protons in the
nucleus of an atom
•
•
•
•
Always a whole number on the periodic table
In a neutral atom, the number of protons and
electrons is equal
The number of electrons can change and the
element remains the same
If the number of protons changes, the element
changes
Distinguishing Between Atoms
Mass Number – total number of protons and
neutrons in an atom
• Measured in atomic mass units (amu) – 1/12 of
the mass of a carbon atom
• Always a decimal number
• Number of neutrons = mass # - atomic #
• Atomic mass is actually the average of an atom
of that particular element
• Isotopes – atoms with the same number of
protons, but a different number of neutrons
Isotopes


Isotopes are atoms of the same element with
different masses.
Isotopes have different numbers of neutrons.
11
C
6
12
C
6
13
C
6
14
C
6
Using the Periodic Table
Atomic #
C
6
Symbol
12.011
Mass #
Periodic Table
Development of the Periodic Table
•
•
Early scientists grouped elements according to
similar properties
Dimitri Mendeleev, a Russian chemist, was the
first to organize the all elements in a systematic
way
o
o
At first there were only 70 elements
Mendeleev even left spaces where he knew would be
discovered
Periodic Table
The Modern Periodic Table
112+ elements on the periodic table
Horizontal rows are called periods.
Vertical columns are called groups or
families
•
•
•
o
o
All elements in a group have similar physical
and or chemical properties because they have
the same number of valence electrons
Valance Electron – electrons in the outermost
energy level of an atom
V. Periodic Table
The Modern Periodic Table
• Arranged by increasing atomic number
•
Periodic Law (periodicity) – pattern on the
periodic table where certain properties
repeat when the elements are arranged
according to atomic number
Periodicity
When one looks at the chemical properties of
elements, one notices a repeating pattern of
reactivities.
METALS
Solids at
room temp
•
High
electrical
conductivity
•
•
Lustrous
•
Ductile
•
Malleable
Non-metals
Various states
of matter at
room temp
•
Poor electrical
conductors
•
•
Non-Lustrous
•
Brittle if solid
Metalloids
Exhibit
properties of
both metals
and nonmetals
•
Periodicity
Main Group Elements ( Group 1, 2, 13 – 18 )
•
•
•
•
Alkali Metals (Group 1) – most highly reactive main
group metals on the periodic table
 Alkali metals are so reactive they are not found in
pure form on earth
Alkaline Earth Metals (Group 2) – exhibit similar
properties to group 1, but are slightly less reactive
Halogens (Group 17) – most highly reactive group of
non-metals
Noble Gases (Group 18) – also called inert gases,
least reactive group of elements on the periodic table
and undergo very few chemical reactions
Periodicity
Transition Metals (Group 3 – 12) – largest group
of elements on the periodic table and contains
about 80% of the metals found on Earth
Inner Transition Metals – bottom two rows of
elements on the periodic table and the least
common elements found on Earth
Periodicity
Inner-transition Metals
• Lanthanide Series – naturally occurring elements
with very similar properties
•
Actinide Series – contain most of the radioactive
elements and all have similar properties
The Periodic Table
INNER-TRANSITION METALS
H
Actinide Series
Noble Gases
Lanthanide Series
Halogens
Alkaline Earth Metals
Alkali Metals
Transition Metals
Chemical Bonding


Bonding occurs between atoms in order to
redistribute their valence electrons in a more
stable way.
The manner of redistribution determines the
type of bonding that takes place
Chemical Bonding

Types of Chemical Bonds
1.
2.
Ionic bonding – chemical bonding resulting
from the electrical attraction between cations
and anions
Covalent bonding – results from the sharing
of valence electrons between two atoms
Chemical Bonding

Factors that effect bonding
1.
Electronegativity – measure of an atoms ability
to attract electrons from another atom in a
chemical bond


Electronegativity is measured from 4.0 to 0, the
higher the number, the higher the electron affinity
or ease of gaining eElectronegativity and the periodic table
 Alkali & Alkaline Earth metals are the least
electronegative
 Non-metals are more electronegative
 Fluorine is the most electronegative
Electronegativity
Electronegativity
Chemical Bonding

Factors that effect bonding
2.
Electron structure





An atom’s electron configuration is the key
factor in determining its electronegativity
Electron-dot notation – visual representation of
an atoms valence electrons
Each atom strives to obtain the same valence
electron structure as a noble gas
Each noble gas has a full valence energy level
because the s & p sublevels are full
Octet rule – atoms combine in such a way as to
completely fill their highest occupied energy
level
Chemical Bonding

Factors that effect bonding
2.
Electron structure





An atom’s electron configuration is the key
factor in determining its electronegativity
Electron-dot notation – visual representation of
an atoms valence electrons
Each atom strives to obtain the same valence
electron structure as a noble gas
Each noble gas has a full valence energy level
because the s & p sublevels are full
Octet rule – atoms combine in such a way as to
completely fill their highest occupied energy
level
Chemical Bonding

Results of chemical bonding
 When
atoms react and combine with one
another they form chemical compounds
 Chemical compounds are represented by
chemical formulas
 Chemical formulas – show the relative numbers
of atoms of each kind in a compound by using
atomic symbols
 Ex:
H2O
CO2
C6H12O6
NaCl
CaBr2
Chemical Bonding

Types of Chemical Compounds
1.
2.
Covalent Compound – group of atoms held
together by the sharing of valence electrons
Ionic Compound – group of atoms held
together through the attraction of oppositely
charged ions
Chemical Bonding



Covalent compounds are always
composed of one or more non-metals
Covalent compounds form when atoms
share their valence electrons
Each atom shares their electrons with the
purpose to have eight valence electrons, or
in accordance with the Octet Rule
Covalent Compounds

Exceptions to the octet rule
 Hydrogen – forms only one bond for a
total of only two valence electrons
 Boron – forms one bond with each of it’s
valance electrons for a total of six
valence electrons
 Other elements, such as Phosphorus,
can have an expanded valence level if
highly electronegative elements are
involved
 Ex: PF5 or SF6
Covalent Compounds

Types of Covalent Compounds
1.
Polar-covalent compound – covalent
compound in which there is an unequal
distribution of electrons

2.
Polar compounds result when an atom in a
compound does not share the electrons equally
across a bond
Nonpolar-covalent compound – covalent
compound where all electrons are shared
equally
Covalent Compounds

Characteristics of Covalent Bonds
 Bond Length – distance needed
between covalently bonded atoms to
minimize their potential energy


The bond length is different depending on
the atoms involved
The bond length is ideal when the distance
between the atoms allows the repulsion of
like charges to equal the attraction of
opposite charges
Covalent Compounds

Characteristics of Covalent Bonds
 Bond Energy – energy needed to break
a covalent bond


The potential energy of the original atoms is
stored in the bonds
Measured in kilojoules per mol (kJ/mol)
Covalent Compounds

Characteristics of Covalent Bonds
 Bond Types


Single Bond – covalent bond in which only
two electrons are shared between atoms
Multiple Bonds – covalent bond in which
multiple pairs of electrons are shared
between atoms
o
o
Double Bond
Triple Bond
Covalent Compounds

Characteristics of Covalent Bonds
 Lewis Structure – using symbols, lines,
and dots to represent the structure and
shape of a molecule

LONE
PAIR
Structural formulas are used to indicate the
basic shape of a molecule without the use of
dots for the lone pairs
O
FREE
e-
Covalent Compounds

Characteristics of Covalent Bonds
 Resonance Structures – refers to
compounds that cannot be represented
following general Lewis structure
procedures


Resonance structures occur when some
elements both share and transfer electrons
to achieve the octet rule
Ex: Ozone ( O3 ) Carbon Monoxide ( CO )
Sulfur Dioxide ( SO2 )
Ionic Compounds




Ionic compounds usually contain at least
one metal
Ionic compounds are composed of atoms
held together due to the attraction of
oppositely charged ions
Ion – any charged particle
Cation – positively charged ion


Created when an atom loses electrons
Anion – negatively charged ion

Created when an atom gains electrons
Ionic Compounds

Characteristics of Ionic Bonds
 Ionic compounds arrange themselves in
a crystal lattice
Ionic Compounds

Characteristics of Ionic Bonds
 Ionic compounds arrange themselves in
a crystal lattice
 Lattice Energy – energy released when
one mole of an ionic crystalline
compound is formed from gaseous ions


Positive energy values indicate energy is
absorbed and stored in the bond
Negative energy values indicate energy is
released
Compound Comparison

Ionic Cmpds





Contains ions
Usually contains a
metal
Strong attraction
between formula
units
High melting and
boiling point
Conduct electricity
when dissolved in
water

Covalent Cmpds
valence e Contains nonmetals
 Weak atraction
between
molecules
 Low melting &
boiling points
 Shares
Polyatomic Ions

Polyatomic ion – charged group of
covalently bonded atoms
 Charge results when there is either too
many or too few electrons to adequately
satisfy the octet rule for each of the
atoms involved
 Polyatomic ions bond with oppositely
charged ions to form part of an ionic
bond
Chemical Formulas
The subscript to the right
of the symbol of an
element tells the number
of atoms of that element
in one molecule of the
compound.
© 2009, Prentice-Hall,
Inc.
Names and Formulas

Chemical formula – indicates
the types and the number of
atoms in a compound



Covalent compound = molecule
Ionic compound = formula units
Chemical formulas serve as a
rough blueprint for a molecule
* First Step *
Determine if the compound
is IONIC or COVALENT
MgBr2
1. Ionic
metals
/
Covalent
non-metals
Names and Formulas

Ionic Compounds
2.
Determine if the compound is
Binary or Tertiary
•
•
Binary compound – composed of only
two different elements
Tertiary compound – composed of
three or more different elements
• All tertiary ionic compounds contain
at least one polyatomic ion
MgBr2
1. Ionic
metals
2. Binary / Tertiary
2 elements
3+ elements
/
Covalent
non-metals
Names and Formulas

Binary Ionic Cmpds w/ Main Group
Metals
3.
Determine if the cation is a Main
Group Metal
•
Group 1, 2 or Aluminum
MgBr2
1. Ionic
metals
/
Covalent
non-metals
2. Binary / Tertiary
2 elements
3+ elements
3. Main Group / Transition Metal
Group 1, 2, Al
Names and Formulas

Binary Ionic Cmpds w/ Main Group
Metals
3.
Determine if the cation is a Main
Group Metal
•
Group 1, 2 or Aluminum
• Name the cation
• Name the anion and change the
ending to -ide
MgBr2
1. Ionic
metals
/
Covalent
non-metals
2. Binary / Tertiary
2 elements
3+ elements
3. Main Group / Transition Metal
Group 1, 2, Al
A. Name cation (metal)
B. Name anion &
change ending to “-ide”
Magnesium Bromide
Li2CO3
1. Ionic
metals
2. Binary / Tertiary
2 elements
3+ elements
/
Covalent
non-metals
Names and Formulas

Tertiary Ionic Cmpds



Tertiary ionic compounds contain
polyatomic ions
Use the chart on p 226 to identify
the polyatomic ion
Use basically the same rules for
naming binary ionic compounds,
but DO NOT change the ending of
the polyatomic ion
Li2CO3
1. Ionic
metals
2. Binary / Tertiary
2 elements
3+ elements
/
Covalent
non-metals
A. Must include polyatomic ion
B. ID the polyatomic ion, use same
rules for binary but DO NOT change the
ending
3. Main Group / Transition Metal
Group 1, 2, Al
A. Name cation (metal)
B. Name anion & change ending to
“-ide” (No change for polyatomic ion)
Lithium Carbonate
Fe2O3
1. Ionic
metals
/
Covalent
non-metals
2. Binary / Tertiary
2 elements
3+ elements
3. Main Group / Transition Metal
Group 1, 2, Al
Any other metal
Names and Formulas

Ionic Cmpds w/ Transition Metals
 Transition metals are different because
the have multiple oxidation numbers
 Oxidation # - charge associated with an
ion in an ionic compound



Determined based on the amount of
electrons lost or gained
Main group elements and polyatomic ions
have constant oxidation numbers
Oxidation numbers must be calculated for
the transition metals according to the anion it
is bonded to

Calculating Oxidation #’s
1.
2.
3.
Determine the ox # for the anion (non-metal
or polyatomic ion)
Calculate the total charge contributed by the
anion
Calculate the unknown ox # for the transition
metal by balancing the charges
Ion
Ox #
X
How many
in cmpd
Total
=
Charge
Names and Formulas

Covalent Compounds





Covalent compounds are composed of
nonmetals bonded together covalently
Covalent compounds do not rely on
charge balance
Naming of covalent compounds uses
prefixes to indicate the number of each
element in the compound
Many covalent compounds have
“names” instead of the stock system of
naming
Ex: C6H12O6 – glucose CH4 - methane
Names and Formulas

Naming Covalent
Compounds
 Name the
compound using
the prefixes
 Mono is
understood if the
first element only
has one atom
 ALWAYS include a
prefix on the
second element
#
Prefix
#
Prefix
1
2
3
4
5
mono
6
7
8
9
10
hexa
di
tri
tetra
penta
hepta
octa
nona
deca
Describing Chemical Reactions

Chemical reaction – process where one or more
substances are changed into one or more different
substances with new chemical properties

Chemical equation – uses symbols, formulas and
molar ratios of compounds to describe a chemical
reaction
6CO2 + 6H2O → C6H12O6 + 6O2
reactants
products
Indications of a Chemical Reaction
1.
2.
3.
4.
Production of light or a transfer of heat/NRG
Production of a gas
Formation of a precipitate (ppt)
Color change
Characteristics of a Chemical Equation
1. The equation must represent known facts
2. The equation must contain correct formulas
for reactants and products
3. Must adhere to the Law of Conservation of
Mass
Matter cannot be created or destroyed
All atoms on the reactant side of the reaction must
equal the atoms on the product side
Balancing Chemical Equations
CH4 + O2 → CO2 + H2O



Make a list of each element on both sides
Start adding COEFFICIENTS to the compounds
to balance the atoms on each side
Balance O’s, H’s, and other single elements last
5 Types of Chemical Reactions
72
1.
Synthesis reaction (composition reaction) – two or
more substances combine to form a new, more
complex compound
2.
A + B → AB
Decomposition reaction – a single compound
undergoes a reaction in which two or more simpler
substances are produced
AX → A + X
NRG
Hagebusch - Chemistry I
4/9/2012
5 Types of Chemical Reactions
73
1.
Single-Displacement reaction (replacement
reaction) – one element replaces a similar element in
a compound
1.
A + BX → AX + B
Double-Displacement reaction – the ions of two
different compounds exchange places in an aqueous
solution to form two new compounds
AX + BY → AY + BX
Hagebusch - Chemistry I
4/9/2012
5 Types of Chemical Reactions
74
5.
Combustion reaction – a substance combines
with oxygen, releasing a large amount of energy
CH4 + O2 → CO2 + H2O
Hagebusch - Chemistry I
4/9/2012
Synthesis Reactions
75

RXNS of Elements with Oxygen & Sulfur



Most all metals will combine with O & S to form
Oxides & Sulfides
Oxygen and Sulfur are both Group 16 elements and
have oxidation #’s of -2
Nonmetals also under go reactions with Oxygen to
produce oxides
2Mg(s) + O2(g) → 2MgO
2H2(g) + O2(g) → 2H2O(g)
Hagebusch - Chemistry I
4/9/2012
Synthesis Reactions (cont.)
76

RXNS of Metals with Halogens


Most metals react with Group 17 elements to form
metallic salts
Fluorine is the most reactive element and will
undergo synthesis with almost every element
2Na(s) + Cl2(g) → 2NaCl(s)
Mg(s) + F2(g) → MgF2(s)
Hagebusch - Chemistry I
4/9/2012
Synthesis Reactions (cont.)
77

RXNS of Metals with Oxides

The oxides of active metals, such as group 1 & 2,
react with water to form metal hydroxides
CaO(s) + H2O(g) → Ca(OH)2(s)
Mg(s) + F2(g) → MgF2(s)

Oxides of nonmetals react with water to form
oxyacids
SO2(g) + H2O(l) → H2SO3(aq)
Hagebusch - Chemistry I
4/9/2012
Decomposition Reactions
78

Decomp. of Binary Compounds


Simplest form of decomposition
Electrolysis – decomposition of a compound by
electric current
H2O(l) → 2H2(g)
ELECTRICITY

+
2O2(g)
Binary oxides decompose through the addition of
heat
Δ
HgO(s) → 2Hg(s) + O2(g)
Hagebusch - Chemistry I
4/9/2012
Decomposition Reactions (cont.)
79

Decomp. of Metal Carbonates

Products are a metal oxide and CO2(g)
Δ
CaCO3(s) →
CaO(s) + CO2(g)

Decomp. of Metal Hydroxides

Products are a metal oxide and H2O(l)
Δ
Ca(OH)2(s) →
CaO(s) + H2O(l)

Decomp. of metal Chlorates

Products are a metallic salt and O2(g)
Δ
2KClO3(s) →
2KCl(s) + 3O2(g)
Hagebusch - Chemistry I
4/9/2012
Decomposition Reactions (cont.)
80

Decomp. of Acids

Certain acids decompose to form non- metal
oxides and H2O(l)
H2CO3(aq) → CO2(g) + H2O(l)
Hagebusch - Chemistry I
4/9/2012
Single-Displacement Reactions
81

Metal Replacement


Occurs when a more reactive metal replaces a less
active metal in a compound
2Al(s) + 3Pb(NO3)2(aq) → 3Pb(s) +
2Al(NO3)3(aq)
Displacement of H+ in H2O by a Metal

Active metals, such as Group 1 react readily with
water to produce metal hydroxides and hydrogen
gas
2Na(s) + 2H2O(l) → 2NaOH(aq)
Hagebusch - Chemistry I
4/9/2012
+
H2(g)
82
Single-Displacement Reactions (cont.)

Displacement of H+ in an Acid by a Metal


Some active metals will replace H+ in acids
Mg(s) + 2HCl(aq) → MgCl2(s) + H2(g)
Displacement of Halogens


One halogen replaces another in a compound
Fluorine is the most reactive
Cl2(g) + 2KBr2(aq) → KCL2(aq)
Hagebusch - Chemistry I
+
4/9/2012
Br2(l)
Double-Displacement Reactions
83

Formation of a Precipitate (ppt)

A precipitate is formed when the cation from one
compound reacts with the anion from another to
form an insoluble compound
2KI(aq) + Pb(NO3)2(aq) → 3PbI2(s) + 2KNO3(aq)

Formation of a Gas
FeS(s) + 2HCl(aq) → 2H2S(g) + FeCl2(aq)
Hagebusch - Chemistry I
4/9/2012
84
Double-Displacement Reactions (cont.)

Neutralization

When strong acids are mixed with strong bases, the
product is a metallic salt and water
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(aq)
Hagebusch - Chemistry I
4/9/2012
Combustion Reactions
85


Spontaneous reaction between a compound
and oxygen gas, producing water vapor
When organic compounds combust, the
products are carbon dioxide and water vapor
2H2(g) + O2(g) → H2O(l)
CH4(g) + O2(g) → 2CO2(g)
Hagebusch - Chemistry I
+
H2O(l)
4/9/2012