Download Electrons in Atoms

CHEMICAL BONDING
Unit 4
LIGHT AND QUANTIZED ENERGY

Nuclear atom and unanswered questions

Scientists found Rutherford’s nuclear atomic model
fundamentally incomplete
Did not explain how electrons are arranged
 Did not address why negatively charged electrons are not
pulled into the atom’s positively charged nucleus


Wave nature of light

Electromagnetic radiation – form of energy that
exhibits wavelike behaviors
Visible light
 Microwaves
 X-rays
 Radio/TV waves

Wavelength (λ) – shortest distance between equivalent
points on a continuous wave
 Frequency (ν) – number of waves that pass a given
point per second

Measured in hertz (Hz)
 Ex: 652 Hz – 652 waves/s

Amplitude – wave’s height from the origin to a crest, or
from the origin to a trough
 All electromagnetic waves travel at a speed of 3.00x108
m/s in a vacuum


Speed of light is represented by “c”
 c = λ ν
Wavelength and frequency are inversely related
 Sunlight passing through a prism is separated into a
continuous spectrum of colors


Particle nature of light

The quantum concept – Max Planck concluded that
matter can gain or lose energy only in small, specific
amounts called quanta
Quantum – minimum amount of energy that can be gained or
lost by an atom
 Equantum = hv
where “E” is energy, “h” is Planck’s constant,
and “v” is velocity
 Planck’s constant = 6.626x10-34 J·s
 Planck’s theory: for a given frequency, v, matter can emit or
absorb energy only in whole-number multiples of hv (1hv, 2hv,
3hv, etc.)
 Analogous to child building a wall with wooden blocks


The photoelectric effect – electrons (photoelectrons) are
emitted from a metal’s surface when light of a certain
frequency shines on the surface
Albert Einstein proposed light has both wavelike and
particlelike characteristics
 Photon – particle of electromagnetic radiation with no mass
that carries a quantum of energy
 Photon’s energy depends on its frequency
 Ephoton = hv


Atomic emission spectra – set of frequencies of the
electromagnetic waves emitted by atoms of the
element
Each element’s atomic emission spectrum is unique and
can be used to determine if that element is part of an
unknown compound
 Neon - light is produced by passing electricity through a
tube filled w/ neon gas


Neon’s atomic emission spectrum consists of several individual
lines of color, not a continuous range of colors as seen in the
visible spectrum
QUANTUM THEORY AND THE ATOM

Bohr model of the atom – Neils Bohr proposed that
elements’ atomic emission spectra are
discontinuous

Energy states of hydrogen
Ground state – lowest allowable energy state of an atom
 Excited state – when an atom gains energy
 The smaller the electron’s orbit the lower the atom’s energy
state


Hydrogen’s line spectrum - when in the excited state the
electron can drop from the higher-energy orbit to a
lower-energy orbit and the atom emits a photon

Fig 5-10

Quantum mechanical model of the atom – Louis de
Broglie accounted for fixed energy levels of Bohr’s
model

Electrons behave as waves
Only half-wavelengths are possible on a guitar b/c the string is
fixed at both ends
 Only whole numbers of wavelengths are allowed in a circular
orbit of fixed radius
 Fig 5-11


Heisenberg Uncertainty Principle – states it is
fundamentally impossible to know precisely both
the velocity and position of a particle at the same
time
Impossible to measure an object w/o disturbing it
 Tried to measure electrons w/ light but b/c a photon has
about the same energy as an electron, the interaction
changes the electron’s position


Quantum mechanical model of the atom – electrons are
treated as waves
Atomic orbital - 3-dimensional region around nucleus that
describes the electron’s probable location
 Fig 5-13


Hydrogen’s atomic orbitals

Principal quantum numbers – indicate the relative sizes
and energies of atomic orbitals

As “n” increases, the orbital becomes larger, the electron
spends more time farther from the nucleus, and the atom’s
energy level increases
Principal energy levels - atom’s major energy levels
(specified as “n”)
 Energy sublevels


Principal energy level 1 consists of a single sublevel; principal
energy level 2 consists of 2 sublevels, etc.
Sublevels are labeled s, p, d, or f according to the shapes of the
atom’s orbitals
 Each orbital may contain at most 2 electrons
 All “s” orbitals are spherical
 All “p” orbitals are dumbbell shaped
 Not all “d” or “f” orbitals have the same shape
 Fig 5-15 and 5-16
 Table 5-2

ELECTRON CONFIGURATIONS

Ground-state electron configurations
Electrons tend to assume the arrangement that gives the
atom the lowest possible energy
 The aufbau principle – states that each electron
occupies the lowest orbital available

All orbitals related to an energy sublevel are of equal energy
 Energy sublevels within a principle energy level have different
energies (Fig 5-17)
 In order of increasing energy, the sequence of energy sublevels
within a principal energy level is s, p, d. and f
 Orbitals related to energy sublevels within one principal
energy level can overlap orbitals related to energy sublevels
within another principal level

The Pauli exclusion principal – states that a maximum
of 2 electrons may occupy a single atomic orbital, but
only if the electrons have opposite spins
 Hund’s rule – states that single electrons w/ the same
spin must occupy each equal-energy orbital before
additional electrons w/ opposite spins can occupy the
same orbitals


Orbital diagrams and electron configuration
notations
Orbital diagram - boxes with zero, one, or two arrows
represent orbitals
 Electron configuration notation – designates the
principal energy level and energy sublevel associated w/
each of the atom’s orbitals and includes a superscript
representing the number of electrons in the orbital


Noble-gas notation

Valence electrons – electrons in the atom’s
outermost orbitals

Electron-dot structure – consists of the element’s
symbol, which represents the atomic nucleus and innerlevel electrons, surrounded by dots representing the
atom’s valence electrons (Lewis dot structure)
FORMING CHEMICAL BONDS
Chemical bond – force that holds 2 atoms together
 Amount of reactivity is directly related to valence
electrons
 Formation of positive ions

Cation – positively charged ion
 Group 1A elements lose 1 valence e-, forming an ion
with a 1+ charge



Ex: By losing an e-, Na acquires the stable outer electron
configuration of Ne
Group 2A elements lose 2 valence e-, forming an ion
with a 2+ charge


Transition metals commonly lose 2 valence e-, forming 2+
ions; however it’s also possible to lose an additional “d”
electron, forming 3+ ion
Formation of negative ions

Anion – negative ion

To designate anions, -ide is added to root of element
 Ex: chloride, sulfide, etc.
Nonmetals form a stable outer electron configuration by
gaining e Group 5A gain 3 e-, forming ions w/ 3- charge
 Group 6A gain 2 e-, forming ions w/ 2- charge

FORMATION AND NATURE OF IONIC BONDS

Formation of an ionic bond

Ionic bond – electrostatic force that holds oppositely
charged particles together


# of e- lost must equal # of e- gained


Oxide – ionic bond between metals and oxygen
Ex: Ca and F form the ionic compound CaF2
Criss-cross method

Properties of ionic compounds
Strong attraction of positive ions and negative ions in an ionic
compound results in a crystal lattice
 Solid ionic compounds are nonconductors of electricity b/c of
the fixed positions of the ions
 Liquid ionic compounds (or those dissolved in water) are
conductors of electricity b/c ions are free to move



Electrolyte – an ionic compound whose aqueous solution conducts an
electric current
Energy and the ionic bond
Endothermic – energy is absorbed during a chemical reaction
 Exothermic – energy is released during a chemical rxn

Formation of ionic compounds from positive and negative ions is
always exothermic
 Attraction of the positive ion for the negative ions close to it forms a
more stable system that is lower in energy than the individual ions
 Lattice energy – energy required to separate one mole of the ions of an
ionic compound
 The more negative the lattice energy, the stronger the force of
attraction
 Directly related to size of the ions bonded



Smaller ions have a more negative value b/c nucleus is closer to, and
thus has more attraction for, the valence e-
Affected by the charge of the ion

Ions w/ larger positive or negative charges have a more negative lattice
energy
NAMES AND FORMULAS FOR IONIC
COMPOUNDS

Formulas for ionic compounds

Formula unit – simplest ratio of the ions represented in
an ionic compound


Ex: KBr – 1:1 ratio; MgCl2 – 1:2 ratio
Determining charge
Binary ionic compound – ionic compound formed by 2 ions
(metal and nonmetal)
 Monatomic ion – one-atom ion
 Ex: Mg2+ and Br Table 8-5 lists common ions of transition metals and groups
3A and 4A
 Oxidation number – charge of a monatomic ion


Compounds that contain polyatomic ions
Polyatomic ions – ions made up of more than one atom
 A polyatomic ion acts as an individual ion
 Exist as a unit, so NEVER change the subscripts of the atoms w/in
the ion
 Table 8-6 lists common polyatomic ions


Naming ions and ionic compounds

Most polyatomic ions are oxyanions – polyatomic ion
composed of an element, usually a nonmetal, bonded to one or
more O atoms

Rules for naming nonmetal-oxyanions (ex: N, S):
Ion w/ more O atoms is named using the root of the nonmetal plus
suffix –ate
 Ion w/ fewer O atoms is named using the root of the nonmetal plus
suffix –ite
 Ex:

NO3-
nitrate
NO2-
nitrite
SO42-
sulfate
SO32-
sulfite

Rules for naming halogen-oxyanions:
Oxyanion w/ most O atoms is named using prefix per-, root of the
nonmetal, and suffix –ate.
 Oxyanion w/ one less O atom is named using root of the nonmetal and
suffix –ate.
 Oxyanion w/ 2 less O atoms is named using root of the nonmetal and
suffix –ite.
 Oxyanion w/ 3 less O atoms is named using prefix hypo-, root of the
nonmetal, and suffix –ite.
 Ex:

ClO4-
perchlorate
ClO3-
chlorate
ClO2-
chlorite
ClO-
hypochlorite

Some things to remember:
Groups 1A and 2A metals have only one oxidation number
 Transition metals and metals on the right side of the periodic table
often have more than one oxidation number
 Ex:

Ions
Formula
Compound Name
Fe2+ & O2-
FeO
Iron(II) oxide
Fe3+ & O2-
Fe2O3
Iron(III) oxide
METALLIC BONDS AND PROPERTIES OF
METALS

Metallic bonds – outer energy levels of the metal
atoms overlap

Electron sea model – all the metal atoms in a metallic
solid contribute their valence electrons to form a “sea”
of electrons


Electrons in the outer energy levels of the bonding metallic
atoms are not held by any specific atom and can move easily
from one atom to the next; these are called delocalized
electrons
Metallic bond – attraction of a metallic cation for
delocalized electrons

Properties of metals
Melting points vary greatly, but in general metals have high melting
and boiling points
 Ex: Hg is liquid at room temp.
0
 Ex: W has a melting point of 3422 C, making it useful for light bulb
filaments and spacecraft parts
 Durable
 Delocalized electrons in metal are free to move, keeping metallic
bonds intact
 Delocalized electrons move heat from one place to another more
quickly than other materials


Metal alloys – mixture of elements that has metallic
properties
Ex: steel is a mixture of iron and at least one other element
 Substitutional alloy – has atoms of the original metallic solid
replaced by other metal atoms of similar size



Ex: sterling silver
Interstital alloy – form when the small holes in a metallic
crystal are filled with smaller atoms

Ex: carbon steel – holes in iron crystal are filled with carbon, making
the solid harder and stronger
THE COVALENT BOND

Covalent bond – chemical bond that results from
the sharing of valence electrons
Generally occurs when elements are relatively close to
each other on the periodic table
 Majority form between nonmetallic elements
 Molecule – formed when 2 or more atoms bond
covalently


Single covalent bonds (sigma bonds)

Group 7A form a single covalent bond


Group 6A share 2 electrons to form 2 covalent bonds


Ex: H2O
Group 5A form 3 covalent bonds with atoms of nonmetals


Ex: F2
Ex: ammonia NH3
Group 4A form 4 covalent bonds

Ex: methane CH4

Multiple covalent bonds
Atoms of C, N, O, and S most often form multiple bonds
 O molecule shares 2 electron pairs, forming a double bond: O2
 N molecule shares 3 electron pairs, forming a triple bond: N2
 Pi bond – parallel orbitals overlap to share electrons


Double covalent bond has 1 sigma and 1 pi bond

Strength of covalent bonds
Bond length – distance between 2 bonding nuclei at the position of
maximum attraction
 As the number of shared electron pairs increases, bond length
decreases



Shorter bond length makes stronger bond
Bond dissociation energy – amount of energy required to break a
specific covalent bond

As 2 atoms are bonded closer together, more bond energy is needed to
separate them
NAMING MOLECULES

Naming binary molecular compounds – prefixes
indicate number of atoms
1-mono
6-hexa
2-di
7-hepta
3-tri
8-octa
4-tetra
9-nona
5-penta
10-deca
Exception – first element never uses mono Common names


Naming acids – if compound produces H+ ions in solution, it
is an acid

Binary acids – contains H and one other element


Use prefix hydro-, root word of other element, suffix –ic, followed by acid
 Ex: HBr is hydrobromic acid
Oxyacids – contains H and a oxyanion
If anion suffix is –ate, it is replaced w/ -ic
 If anion suffix is –ite, it is replaced w/ -ous
 Ex:
HNO3 is nitric acid
HNO2 is nitrous acid


Considered acids only when in solution (i.e., dissolved in water)

Ex: HCl at STP is hydrogen chloride gas but in solution is hydrochloric acid
MOLECULAR STRUCTURES

Structural formulas – letter symbols and bonds
show relative positions of atoms
H is always on the end
 Central atom is usually closer to the left on the periodic
table
 If central atom is not surrounded by 4 e- pairs, it does
not have an octet

Must have double or triple bond
 Ex. problems 9-3 – 9-5



Resonance structures – more than one valid Lewis structure can
be written; differ only in position of e- pairs and never atom
positions
Exceptions to Octet Rule
1.
Uneven number of valence electrons - some molecules have an odd
number of valence electrons and can’t form an octet

2.
Suboctet - some compounds form w/ fewer than 8 electrons present



3.
Ex: NO2
Reactive and can share an entire pair of electrons donated by another
atom
Coordinate covalent bond – when one atom donates a pair of electrons to
be shared w/ an atom or ion that needs 2 electrons to become stable
Ex: BH3
Expanded octet - some compounds have central atoms that contain
more than 8 valence electrons

Ex: PCl5
MOLECULAR SHAPE

Valence Shell Electron Pair Repulsion (VSEPR)
model – used to determine molecular shape





Based on arrangement that minimizes repulsion of
shared and unshared pairs of electrons around the
central atom
Bond angle – angle formed by any 2 terminal atoms and
the central atom
Lone pairs of electrons are not shared between 2
nuclei, therefore they occupy a slightly larger orbital
than shared electrons
Shared bonding orbitals are pushed together
slightly by lone pairs, therefore they occupy a slightly
smaller orbital than lone pairs
Table 9-3

Hybridization – atomic orbitals are mixed to form new,
identical hybrid orbitals


Hybrid orbitals – contain one electron that it can share w/ another
atom
Ex: methane CH4
ELECTRONEGATIVITY AND POLARITY

Polar covalent bond – unequal sharing of electrons
due to differences in electronegativity


Ex: water H2O
Microwave ovens
Water molecules are constantly moving as they align
themselves w/ the changing polarity of the microwave
energy field
 How it Works, p. 270
