Chapter 6. chemical bond an electromagnetic attraction between ______ ____ ______ __that binds atoms...

Chapter 6

chemical bond• an electromagnetic attraction between

______ ____ ______ __that binds atoms together

• ionic bonding – electromagnetic attraction between atoms that ______ _

- attraction between ions of opposite charge

• ______ ________– electromagnetic attraction between atoms that share e-

• covalent bonds form between different _________• each nonmetal needs _____________to be

electrically stable(full outer NRG level)• nonmetals ________until outer NRG level is

full• example – hydrogen & chlorine

• example – hydrogen & oxygen

• ionic bonds form between ______ ___ ________• metals have extra valence e- • nonmetals need more e-

• example – NaCl• Na goes thru ________• Na Na+ + e-

• Cl goes thru _________• Cl + e- Cl-

• opposite charged ions attract together• Na+ + Cl- Na+Cl-

Ionic or Covalent????• determined by electronegativity

differences(pg 161)

• electroneg diff > 1.7 ______ ______• greater than 50% ionic character

• electroneg. diff < 1.7 ________ _____• less than 50% ionic character

• two possibilities if covalent1) atoms have different electroneg.

• e- sharing is _________• creates ________ charges on the atoms

• not as charged as an ion• δ+ or δ- (lower case Greek letter delta) = partial charge• δ+ = _______ ________, δ- = _______ ________

• partial charges on atoms = _______• polar covalent bond = ________ sharing of e-

2) atoms have similar electroneg. • same element or elements close

together on periodic table• e- sharing is ______• no charges or poles created• no poles = _________ covalent bond

• ________ covalent bond• formation of _______(δ+ or δ-) charges• covalent bonds with more than 5% ionic

character and less than 50% ionic character• ____ > electronegativity difference > ___

• nonpolar covalent bond • no ___________ formed• aka. ______ covalent bond• covalent bond with less than 5% ionic

character• electronegativity < _____

If the following elements were to bond with sulfur what type of bond would be formed?

hydrogen

cesium

chlorine

2.5 - 2.1 = 0.4

_____ ______

_____

_____ ______

2.5 - 0.7= 1.8

3.0 - 2.5 = 0.5

6.2 Covalent bonding & molecules________ – stable, independent particle

created when atoms share e- • nonmetals will share e- to achieve a ___

______ ______ _______

• covalent bond• co = _______• valent = _______ __

• hydrogen atoms• 1 valence e- H-1s1 • needs to ____ ____ ___ __ for full outer

NRG level• no metals present = no oxidation• 2nd H atom also would like to gain e-

• creates a _________ ________• the most probable location of shared e-

around both nuclei

• ________ _______= covalent bond• each atom has e- config of He-1s2 • both atoms have a ___ ____ ____ ____

• 1 e- of their own, 1 e- they are sharing

• NRG is released when bond forms• ____ ______ together than apart

• same as ionic bond

• the more NRG released = _______ bond length• bond length – the avg _______ between two

bonded atoms• bond NRG – NRG required to _______ a

covalent bond• shorter bond length = greater bond NRG

• _______ _______– shows actual number of atoms in stable independent particles• H2 , CCl4 , C6H12O6

• does not need to be empirical

• _______ ____formula – shows atoms and valence electrons, shared pairs between symbols• H:H

• _________ formula – shows atoms and covalent bonds represented as dashes• H-H

• _________ elements – molecules created when only two atoms of the same element bond together• H2, F2, Cl2, Br2, I2

• H-H, F-F, Cl-Cl, Br-Br, I-I• all share __ ______ of e-

• single covalent bond

• O2 & N2 form ________ covalent bonds

• O – 1s22s22p4 (needs 2 more e- )• shares __ _____ of e-

• double covalent bond

• N - 1s22s22p3 (needs 3 more e- )• shares __ ______ of e-

• triple covalent bond

• _______ dot structures• method used to create model of

molecules• give the basic _______/______ of a

molecule1) total valence e- of all atoms in molecule

• HCl = 8 valence e-

2) write symbols for each element• least number of = ________ atom• hydrogen is always ________ atom• H Cl

3) add a _____ _____ __ between atoms bonding together

• H : Cl valence e- remaining = 6

4) add remaining e- in pairs to exterior atoms to form _______

• if e- remain add them to interior atoms to form octets

• not H, H forms duets

• if necessary, move e- pairs to create octets

• H : Cl

5) all _______ pair of e- become dashes

• H - Cl

::

:

:

::

6) follow ________ for shape• VSEPR – valence shell _______ pair

repulsion theory• model used to predict ______/_______of a

molecule based on the repulsion of e- pairs• e- pairs repel each other to maximum

distance• in 2-dimension = 90o

• in 3-dimension = 109.5o

• ________ – ability to create two distinct models that accurately represent the structure of a molecule

6.5 Molecular Geometry• 3-dimensional shapes determine the

_________ and ________ properties of molecules• example – sucrose- its molecular shape fits

the nerve receptors of the tongue for sweetness• sugar substitutes(Splenda, Nutrasweet, …)

have similar ________ as sucrose

1) ________• all atoms lie in a straight line• HCl, CO2

2) _____• atoms do not lie in straight line• e- pairs point to 4 corners of _________

3) _______ _________• 3 atoms bonded to central atom and a

pair of nonbonding e-

4) ________ _________• 3 groupings of e- around central atom• all atoms lie in same plane

5) ___________• 4 groupings of e- around central atom

• ___________• formation of a new hybrid atomic orbital from

2 or more atomic orbitals of _______ _______• example - carbon

• ____ hybridization

• _____ hybridization

• _____ hybridization

• ________ molecules• a molecule that has oppositely, partially

charged atoms on opposite sides• aka - _______

• a molecule that has an __________ distribution of charge

• partial charges not evenly distributed around ________ atom

• polar molecules:• must have polar bonds• partial charges are unevenly distributed

• HCl• bond is _______(electroneg diff = ___)• H = ___, Cl = ___• _____ shape• _________ molecule

• N2

• bond is ________(electroneg diff = ___)• ________ molecule

• H2O• bonds are ______(electroneg diff = ___)• H = ___, O = ____• _____ shape• _______ molecule

• NH3

• bonds are _____(electroneg diff = ___)• H = ___, N = ____• _________ ________shape• ________ molecule

• CCl4• bonds are ______(electroneg diff = ___)• C = ____, Cl = _____• ___________ shape• __________ molecule

Intermolecular Forces

• Intermolecular Force – the force of attraction between individual __________

• no charges

– causes solids and liquids to stick together• the “___________” that holds the molecules together in the liquid

and solid phases

– _____molecular forces are much weaker than ______molecular forces.• _______molecular forces – covalent bonds that exist between

the atoms bonded in a molecule

____________ force

____________ force/ covalent bond

– When a substance ______ or __________, the intermolecular forces are broken (not the covalent bonds).• the “velcro” pulls apart, not the molecule.

Types of Intermolecular Forces1. __________________ Forces

- force of attraction between two polar molecules

2. _____________ Bonding

- force of attraction between molecules that have hydrogen and a very electronegative element(F,O,N)

3. ___________________Forces

- force of attraction between all molecules(polar and nonpolar molecules) and neutral atoms

• Dipole – Dipole force– caused by ________ charges created thru

unequal sharing of electrons in a polar bond

• Hydrogen Bond– force of attraction between molecules that have

_________ and a very _________________ element(____, ____, ____)• Special type of dipole-dipole forces.• stronger than _______________.

– ~ 5 to 10% the strength of a covalent bond. – dipole-dipole force is ~1% strength of a covalent bond

• unique properties of _______ attributed to hydrogen bonding– unusually high __________ _________

• strong adhesion and cohesion

– high ________ ____ ______ _____compared to similar molecules• H2O – boils = 100oC, melts = 0oC

• H2S – boils = -60oC, melts = -85.5oC

• H2Se – boils = -41.3oC , melts = -65.7oC

• H2Te – boils = -2oC, melts = -49oC

– unusual ____________ phenomena• most dense @ 4oC

– one of only a few substances(Ga, Bi, Si) in which solid is less dense than liquid

» ______ sinks in own liquid• ice is 89% as dense as water

• above 4oC – thermal expansion/contraction

• below 4oC – hydrogen bonding causes molecules to spread

out– thermal contraction is countered by hydrogen

bonds

• when _________, molecules organize and push further apart(less dense) @ _____ – hydrogen bonds reach maximum rigidity

• when _______, molecules disorganize and crowd closer together(more dense) @ 0oC – hydrogen bonds lose rigidity and become

flexible

• examples of hydrogen bonding– Hydrogen bonding between guanine and cytosine, and adenine

and thymine in DNA

• examples of hydrogen bonding– the strength and shape of hair

• examples of hydrogen bonding– help enzymes bind to their substrate– help antibodies bind to their antigen – aids with the hardening of concrete– life would be very different/nonexistent on earth if

hydrogen bonding didn’t exist• “It is found that if the hydrogen bond strength was slightly

different from its natural value then there may be considerable consequences for life. At the extremes water would not be liquid on the surface of Earth at its average temperature if the hydrogen bonds were 7% stronger or 29% weaker. The temperature of maximum density naturally occurring at about 4°C would disappear if the hydrogen bonds were just 2% weaker. Major consequences for life are found if the hydrogen bonds did not have their natural strength. Even very slight strengthening of the hydrogen bonds may have substantial effects on normal metabolism.”

• Water’s Hydrogen Bond Strength by Martin Chaplin (http://arxiv.org/ftp/arxiv/papers/0706/0706.1355.pdf)

• London Dispersion force(aka – van der Waals force)

– ____________ of all intermolecular forces– found in ____ molecules and ________ atoms

– two adjacent neutral molecules/atoms can affect each other

– the nucleus of one molecule/atom attracts the electrons of the adjacent molecule/atom

– for an _________, the electron clouds become distorted.

– creates a ___________ dipole(called an _____________ dipole)

• One instantaneous dipole can ________ another instantaneous dipole in an adjacent molecule (or atom).– induced dipoles – ________ dipoles created

by adjacent instantaneous dipoles– the forces between instantaneous dipoles are

called ________ _________ ______.

• Strength of London dispersion forces: – related to ___________ _________

• for similar substances• reiterated by boiling points*

– *for atoms and molecules that only have induced dipole-dipole intermolecular forces

Nobel Gas (g/mole) boil pt. (K)

He 4.0 4.6

Ne 20.2 27.3

Ar 39.9 87.5

Kr 83.8 120.9

6.3 Ionic Bonding_________compounds

• compound formed when metals _______ and __________ reduce

• Ca Ca2+ + 2e- • Cl + e- Cl-

• Ca2+ + 2Cl- CaCl2

• formula unit• the _________ ratio of ions in an ionic

compound

2 2

• Ionic compounds– ions are held together by ______ ______

• structure of an ionic compound is an ionic crystal• each ion is bonded to all the oppositely charged

ions around it

• ionic bonds are _________ bonds– size of ion(smaller = stronger)– charge of ions(more = stronger)

• ionic compounds have high melting and boiling points

Salt melt pt boil pt

KCl 770oC sublimes @ 1500oC

- chemically decomposes

NaCl 801oC 1413oC

CaF2 1423oC 2500oC

– high melting and boiling points indicate _______ ________

• metallic bonds• force of attraction between _______ ____• caused by loosely held ________ ___• cations surrounded by a pool of e-

• good conductor of _________• good conductor of _________• malleable, ductile, tenacious