Chapter 4
Atoms and Chemical Bonding
Continuous Spectra
Roy G. Biv
RedOrangeYellowGreenBlueIndigoViolet
Line Spectrum
Model of an Atom
• Must explain many things, one of which is the line spectra
• First to come up with a decent model was Bohr• Visualized the electrons in an atom orbiting around
the nucleus like a planet around a sun• Each orbit was associated with a definite energy
(no in-between levels)• Took a “quantum” of energy (fixed amount) to
move from one level to another
Electron Configuration of the AtomBohr (1885-1962)
Electron Energy Levels
Bohr’s Model
• Energy levels each can hold a different (maximum) number of electrons
• Energy level (n) = 1, 2, 3, 4, …• # electrons (2n2)= 2, 8, 18, 32, …
• Example: Na has 11 electrons
– 2e – 8e – 1e (1st) (2nd) (3rd) -> has 17 empty spots in 3rd level
Bohr’s Model
• Worked well for hydrogen• Worked okay for about the next 20 atoms• Didn’t work well at all for anything past the transition
metals• Need new model…
Wave Mechanical Model of the AtomSchrodinger (1881-1961)
Retains Bohr’s energy level concept Distinguishes orbitals at each energy level Orbitals identified as: s, p, d, f
Orbitals in the First Four Energy Levels in the Wave-Mechanical Model of the Atom
Building Atoms with Wave-Mechanical Model
Electrons are added starting at the first level Superscripts are used to indicate the number of
electrons in each orbital
Electron Arrangements of
the First 20 Elements in the Periodic Table
Energy Levels of Oribitals
An Aufbau Diagram
Using the Periodic Table to Determine Electron Configuration
The Shape of an s Orbital
The Shape of a p Orbital
Chemical Properties and the Periodic Table
Chemical properties and electron configuration correlate Alkali metals all have one s electron in their highest
energy level
Li 1s2 2s1
Na 1s22s22p63s1
K 1s22s22p63s23p64s1
Rb 1s22s22p63s23p64s23d104p65s1
Valence ElectronsLewis (1916)
Valence electrons are the electrons in the atom’s highest numbered energy level.
Octet Rule
• Atoms tend to gain or lose electrons to have eight valence electrons• Same as noble gases• He and H are exceptions, get only two valence
electrons
Stable Electron Configurations
• Valence electrons – outermost level with electrons• Core electrons – all other electrons in an atom• Isoelectronic – same number of valence electrons
• Example: O2-, F-, and Ne all have 18 e-’s
Electron-Dot Structures
• Valence electrons represented by dots• Electron-dot symbols
– Examples: Na•, •Mg•, …
Lewis Dot Structures
• The Lewis dot representation (or Lewis dot formulas)• convenient bookkeeping method for valence
electrons• electrons that are transferred or involved in
chemical bonding
Chemical Bonds
• Forces responsible for holding together atoms in molecules and ions in crystals
• Determine shape of molecules• Predict chemical and physical properties of materials• Related to arrangement of electrons in compounds• The electrons involved in bonding are usually those in
the outermost (valence) shell.
Chemical Bonding
• Chemical bonds are classified into two types:• Ionic bonding results from electrostatic attractions
among ions, which are formed by the transfer of one or more electrons from one atom to another.
• Covalent bonding results from sharing one or more electron pairs between two atoms.
Ionic Bonding
• Remember• cations or positive (+) ions
• atoms have lost 1 or more electrons• anions or negative (-) ions
• atoms have gained 1 or more electrons
Ionic Bonding
Atoms lose or gain electrons to form ions Cations are positive ions Anions are negative ions
Ionic compounds are held together by electrostatics- the positive charge of the cation attracting the negative charge of the anion.
Ionic Bonding Continued
Ionic Bonds
• Na+ and Cl–
• Opposite charges attract • Ions organize themselves in orderly manner
• Crystal of NaCl
Structures of Ionic Compounds
• extended three dimensional arrays of oppositely charged ions
• high melting points because coulomb force is strong
Ionic Bonding
• We can also use Lewis formulas to represent the neutral atoms and the ions they form.
Li + F...
.... .
Li+
F[ ]...... ..
Naming Ions
• For cations, simple positive ions• Add the word ion• Examples: Na+ – sodium ion
Al3+ – aluminum ion• For anions, simple negative ions
• Change the usual ending to -ide• Examples: Cl– – chloride
S2– – sulfide
Ionic Compound
Ionic compounds are formed primarily when metals on the left side of the periodic table react with nonmetals on the right side of the periodic table.
Transition metals also form ionic compound Their behavior is less predictable Iron forms Fe2+ or Fe3+
Copper forms Cu+ or Cu2+
Naming Binary Ionic Compounds
• Two components in compound
Common Ions and Their Position in the Periodic Table
Polyatomic Ions
Polyatomic means “many-atom” ion
Mem
orize These
Polyatomic Ions
• “Many-atom” Ions
• Example: NH4+, OH-, CH-
• Covalently bonded groups of atoms, that tend to stay together
Na + O + H O H -- + Na+
Equations
+
+
+
-
-
-
Polyatomic Ions
• Charged groups of atoms that remain together through most chemical reactions
Covalent Bonds
• Bond formed by a shared pair of electrons• Gives atom an octet of electrons
• Shared pair of electrons – bonding pair• Other electrons not involved in bonding –
nonbonding pairs
Covalent Bonding
• covalent bonds formed when atoms share electrons• share 2 electrons - single covalent bond• share 4 electrons - double covalent bond• share 6 electrons - triple covalent bond• attraction is electrostatic in nature
• lower potential energy when bound
Covalent Bonding
extremes in bonding:• pure covalent bonds
• electrons equally shared by the atoms• pure ionic bonds
• electrons are completely lost or gained by one of the atoms
• most compounds fall somewhere between these two extremes
Covalent Bonds
How do two identical atoms bond to form molecules such as H2, N2, or O2?
Neither atom is more likely than the other to transfer an electron
The two atoms have to share electrons
Covalent Bonding
Number of Covalent Bonds/Element
• Follow the electron-dot rules for the following elements
Covalent Bonding
• Lewis dot representation • H molecule formation
+H. H . H H.. or H2
Covalent Bonding
• Lewis dot representation • H molecule formation
• HCl molecule formation
+H. H . H H.. or H2
H Cl H Cl+...
.... ..
..
..
... or HCl
Lewis Dot Structures
• homonuclear diatomic molecules
• hydrogen, H2
• fluorine, F2
• nitrogen, N2
H HorH H..
F F.. .. ....
..
.. ..F F
.. .... ..
.. ..or
N N········ ·· N N·· ··or
Lewis Dot Structures
• heteronuclear diatomic molecules• hydrogen halides
• hydrogen fluoride, HF
• hydrogen chloride, HCl
• hydrogen bromide, HBr
or ··H F··
··H F..
······
or ··H Cl··
··H Cl..
······
or ··H Br··
··H Br..
······
Lewis Dot Structures
• water, H2O
H
H
O··
····
··
Lewis Dot Structures
• ammonia molecule , NH3
H
H
N··
····
·· H
Lewis Dot Structures
• ammonium ion , NH4+
H
H
N··
····
·· H
H +
N - A = S rule
• N = # of electrons needed to be noble gas• 8 for everything except H or He • Only 2 for H or He
• A = # of electrons available in outer shells• equal to group #• 8 for noble gases
N - A = S rule
• for ions• add one e- for each negative charge• subtract one e- for each positive charge
• central atom in a molecule or polyatomic ion is determined by:• atom that requires largest number of
electrons• for two atoms in same group - less
electronegative element is central
Drawing Lewis Dot Formulas
• Example: Write Lewis dot and dash formulas for hydrogen cyanide, HCN.
• N = 2 + 8 + 8 = 18• A = 1 + 4 + 5 = 10• S = 8
Drawing Lewis Dot Formulas
• Example: Write Lewis dot and dash formulas for hydrogen cyanide, HCN.
• N = 2 + 8 + 8 = 18• A = 1 + 4 + 5 = 10• S = 8
H C N·· ·· ···· H C N ··or··
Drawing Lewis Dot Formulas
• Example: Write Lewis dot and dash formulas for the sulfite ion, SO3
2-.
• N = 8 + 3(8) = 32• A = 6 + 3(6) + 2 = 26• S = 6
Drawing Lewis Dot Formulas
• Example: Write Lewis dot and dash formulas for the sulfite ion, SO3
2-.
• N = 8 + 3(8) = 32• A = 6 + 3(6) + 2 = 26• S = 6
O S O
O··
····
····
··
··
··
····
··
····
2-O S
O
O·· ·· ··
······ ··
······
2-or
Naming Covalent Compounds
Use prefixes to indicate the number of each kind of atom
Examples:
Carbon Monoxide
Carbon Dioxide
Trinitrogen Pentoxide
Polar and Nonpolar Covalent Bonds
• nonpolar covalent bonds• electrons are shared equally• symmetrical charge distribution
• must be the same electronegativity to share exactly equally (typically by being the same element)
• H2
• N2
H HorH H..
N N········ ·· N N·· ··or
Polar and Nonpolar Covalent Bonds
• polar covalent bonds• unequally shared electrons• assymmetrical charge distribution• different electronegativities
ElectronegativityPauling (1901-1994)
Electronegativity is the relative tendency of an atom in a molecule to attract a shared pair of electrons in a bond to itself.
The most electronegative element is fluorine and it is given a value of 4.0.
The higher the electronegativity value of an atom, the greater is the ability of an atom of that element to attract electrons to itself.
Electronegativity Values for the Representative Elements
Polar Molecules
When hydrogen and chlorine react to form HCl, a polar molecule is formed.
Continuous Range of Bonding Types
• all bonds have some ionic and some covalent character• HI is about 17% ionic
• greater the electronegativity differences the more polar the bond
Polar Covalent Bonds
• If elements do not have the same electronegativity, they get unequal sharing of electrons
Polar and Nonpolar Covalent Bonds
bondpolar very 1.9 Difference
4.0 2.1 ativitiesElectroneg
F H
1.9
Polar and Nonpolar Covalent Bonds
• Electron density map
of HF• blue areas – low
electron density• red areas – high
electron density• Polar molecules have separation of centers of
negative and positive charge
Polar and Nonpolar Covalent Bonds
bondpolar slightly 0.4 Difference
2.5 2.1 ativitiesElectroneg
I H
0.4
Polar and Nonpolar Covalent Bonds
• Electron density map
of HI• blue areas – low
electron density• red areas – high
electron density• Notice that the charge separation is not as big as for
HF• HI is only slightly polar
Polar Molecules
• Molecule with an overall partial charge• Can have polar bonds and be non-polar
molecule
• O = C = O O
H HNo overall charge - Symmetrical
Overall charge - Asymmetrical
Polar Molecules Continued
Carbon Dioxide, CO2 has polar bonds but because of its symmetrical shape it is a nonpolar molecule.
O=C=O
Intermolecular Forces
• Glue that holds matter together• Melting and boiling points measure the relative
strength • Ionic forces – strongest
• Found in salts• NaCl melts at 800°C
Hydrogen Bonds
Hydrogen Bonding
• Hydrogen must be attached to electronegative atom• N, O, F
• Plays important role in biological systems
Hydrogen Bonding
• Particularly strong dipole-dipole interaction• Occurs between Hydrogen and:
• Oxygen, Nitrogen and Flourine
• Why? High electronegativity of O, N, F
Dipole Moments
• in molecules• some nonpolar molecules have polar bonds
• 2 conditions must hold for a molecule to be polar
units Debye 0.38 units Debye 1.91
I-H F-H
--
Dipole Moments
• There must be at least one polar bond present or one lone pair of electrons.
• The polar bonds, if there are more than one, and lone pairs must be arranged so that their dipole moments do not cancel one another.
Dipole Forces
• Not as strong as ionic forces• Must have polar molecule
• HCl melts at –112°C
Dipole-Dipole Interaction
• Caused by permanent dipoles (partial charges) on the molecule
Cl ClH H
Cl H Cl H- + - +
+ +--
London Forces
Dispersion Forces
• Present in all molecules• Weak momentary attractive forces
• Arise for electrons moving about in molecules and atoms
• Strong in larger molecules• Important in nonpolar compounds
HW Suggestion
• Ch 4:• 3, 4, 6, 9, 11, 13, 18, 19, 20, 23, 24, 27, 29,
33, 34, 35, 37, 39, 40, 43, 45, 46, 49, 53, 55
• Wednesday: Lab! Bring finished prelab writeup, wear closed-toe shoes, hair up, bring goggles, etc…!
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