Chapter 8 & 9 Flashcards by Jenna B

Electrons are transferred; electronegativity differences are generally greater than 1.7; formation is always exothermic

ionic bonding

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Electronegativity difference is not the final determination of __ __; compounds are ionic if they conduct electricity in their __ state

ionic charactermolten

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__ are larger than __ because of higher effective nuclear charge

anions, cations

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Ionic compounds form solid crystals at __ temperatures; ionic compounds organize in a characteristic __ __ of alternating positive and negative ions; all __ are ionic compounds and form crystals

ordinarycrystal latticesalts

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Properties of Ionic compounds:Structure:Melting point:Boiling point:Electrical conductivity:Solubility in H2O:

Structure: crystalline solidsMelting point: generally highBoiling point: generally highElectrical conductivity: excellent conductors in molten and aqueous statesSolubility in H2O: generally soluble

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Ionic size from largest to smallest with charges:3+, 2+, +, -, 2-

2- > - > + > 2+ > 3+

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Greater the effective nuclear charge, __ the ionic size

smaller

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nonmetal atoms come together to form bonds to become more stable by sharing electrons to achieve a noble gas electron configuration; type of bond that holds atoms together in a molecule

covalent bonding

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atoms form covalent bonds in such a way that all atoms get 8 electrons in its highest occupied energy level, except hydrogen which only needs 2 electrons

octet rule

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ability of an atom in a bond to attract shared electrons to itself; __ across a period and up a group

electronegativityincreases

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electrons are equally shared; electronegativity difference of 0 - .4

Non polar-covalent bonds

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electrons are unequally shared; electronegativity difference of .5 - 1.7; has two poles- a positive and negative end

polar covalent bond

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Polar bonds:electrons concentrate around the more __ atom in a molecule; atom gains a partial __ charge; since electrons spend __ time around the other atom, other atom gains a partial __ charge

electronegativenegativelesspositive

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Covalent bonding forces:Electron - electron __ forcesProton - proton __ forcesElectron - proton __ forces

repulsiverepulsiveattractive

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1) Longest and weakest bonds2) shortest and stronger than single bonds3) shortest and strongest type of covalent bond

1) single bonds2) double bonds3) triple bonds

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distance between the nuclei of two covalently bonded atoms; optimum distance __ attraction, __ repulsions

bond lengthmaximizingminimizing

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Shorter bond length = __ bond energy = __ bondBonds b/w elements become __ and __ as multiplicity increases

higherstrongershorterstronger

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polarity depends on __ __ and __ around central atom

individual bondsgeometry

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representative elements generally form __ to achieve a noble gas configuration

ions

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ions from representative metals are usually __ with one of the noble gases (have the same electron configuration)

isoelectronic

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Transition metals lose __ electrons before the __ electrons when forming cations

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distance from the nucleus to the outermost electrons in an ion

ionic radii

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An atom loses electrons to form a __; it has a __ radius than its corresponding atom

cationsmaller

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An atom gains electrons to form an __; it has a __ radius than its corresponding atom

anionlarger

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innermost electrons belonging to filled electron shells

core electrons

electrons in the outermost shell; they combine with other atoms with unfilled shells to become stable; lead to chemical bonds and reactions b/w atoms

valence electrons

Many bonds forming, a lot of __ is released; __ released more than makes up for the __ __ required

energyenergyionization energy

energy required to convert and ionic compound into gaseous ions; quantitative measure of the strength of the ionic bonds in the compound

lattice energy

Relative strength of an ionic bonds is determined by the following:1) Charges of ions: __ the charge, the stronger the bond2) Distance b/w two ions: __ distance, stronger the bond

highershorter

The strength of the ionic bond is generally determined foremost by the __, and only if the __ are similar does one compare the distance b/w nuclei to determine the strength of the bond and the resulting lattice energy

chargescharges

Estimate heat of formation for NaCl:Na(s) + 1/2 Cl2(g) -> NaCl(s)Lattice energy = -786 kJ/molIonization energy for Na = 495 kJ/molElectron affinity for Cl = - 349 kJ/molBond energy of Cl2 = 239 kJ/molEnthalpy of sublimation for Na = 109 kJ/mol

Na(s) -> Na(g) = 109 kJ/molNa(g) -> Na+(g) + e- = 495 kJ/mol1/2Cl2(g) -> Cl(g) = 1/2(239 kJ/mol)Cl(g) + e- -> Cl-(g) = - 349 kJ/molNa+(g) + Cl-(g) -> NaCl(s) = -786 kJ/molTOTAL:Na(s) + 1/2Cl2(g) -> NaCl(s) = -412 kJ/mol

ions form in a __ state

gaseous

energy required to break a particular bond in the gas phase; always positive since it takes energy to break a bond; quantitative measure of a bond's strength (__)

bond energystability

Enthalpy of formation = sum of bonds broken (energy __) - sum of bonds formed (energy __)

requiredreleased

__ bonds always requires energy (endothermic, +)__ bonds always releases energy (exothermic, -)

breakingforming

covalent compounds involve atoms of __ only

nonmetals

the term __ is used exclusively for covalent bonding

molecule

When drawing lewis structure, put least __ atom as central atom and surround with other atoms

electronegative

ternary oxyacids contain __, __, and one other element

oxygenhydrogen

In ternary oxyacids, the central atom is the __ element which is surrounded by __ and the __ atoms are bonded directly to the __ atoms

otheroxygenhydrogenoxygen

what atoms are satisfied with less than an octet?

Be, and B

Atoms is and beyond the __ period can have more than 8 electrons when in a compound

third

hypothetical charge an atom would have if bonding electrons are shared equally and lone pairs belong solely to a single atom

formal charge

formal charge = total # of __ electrons - total # of __ electrons - total # of __

valencenonbondingbonds

For neutral molecules, sum of formal charges must equal _. For ions, the sum of formal charges must equal __

0charge

Lewis formulas with no formal charges is __ to one with formal charges; Lewis formula with large formal charges is __ plausible than one with lower formal charges; most plausible formula has negative formal charge on __ electronegative atom

preferablelessmore

electrons spread between more than two atoms

delocalized electrons

one of two or more lewis structures representing a single molecule that cannot be described fully with only one lewis structure

resonance structure

the 3-D arrangement of atoms in molecules; responsible for many physical and chemical properties of molecules

molecular geometry (shape)

a molecule's shape results from the electrons on its central atom orienting themselves to be as far away from each other as possible

VSEPR model

If there are only two atoms, the molecule must be __

linear

If there are more than two atoms in the molecule, the shape depends on the number of __ around the central atom

electrons

Hydrogen forms __ covalent bond

one

Oxygen and Sulfur form __ covalent bonds (__ double or __ single bonds)

212

Nitrogen and Phosphorus form __ covalent bonds(__ triple bond, __ single bonds, or __ double & __ single bond)

3, 1, 3, 1, 1

Carbon and Silicon form __ covalent bonds (__ double bonds, __ single bonds, __ triple & __ single, or __ double & __ single bonds)

4, 2, 4, 1, 1, 1, 2

Resonance bonds are __ and __ than single bonds. They are also __ and __ than double bonds

shorterstrongerlongerweaker

electron pairs can be thought of as belonging to pairs of atoms when bonding

LEM model

Shape: Shared electrons (bonds): 2Lone pairs of electrons: 0Hybridization:Bond angle(s):

Shape: linearShared electrons (bonds): 2Lone pairs of electrons: 0Hybridization: spBond angle(s): 180

Shape:Shared electrons (bonds): 3Lone pairs of electrons: 0Hybridization:Bond angle(s):

Shape: trigonal planarShared electrons (bonds): 3Lone pairs of electrons: 0Hybridization: sp2Bond angle(s): 120

Shape:Shared electrons (bonds): 2Lone pairs of electrons: 1Hybridization:Bond angle(s):

Shape: bentShared electrons (bonds): 2Lone pairs of electrons: 1Hybridization: sp2Bond angle(s): 120>

Shape:Shared electrons (bonds): 4Lone pairs of electrons: 0Hybridization:Bond angle(s):

Shape: tetrahedralShared electrons (bonds): 4Lone pairs of electrons: 0Hybridization: sp3Bond angle(s): 109.5

Shape:Shared electrons (bonds): 3Lone pairs of electrons: 1Hybridization:Bond angle(s):

Shape: trigonal pyramidalShared electrons (bonds): 3Lone pairs of electrons: 1Hybridization: sp3Bond angle(s): 109.5>

Shape:Shared electrons (bonds): 2Lone pairs of electrons: 2Hybridization:Bond angle(s):

Shape: bentShared electrons (bonds): 2Lone pairs of electrons: 2Hybridization: sp3Bond angle(s): 109.5>

Shape:Shared electrons (bonds): 5Lone pairs of electrons: 0Hybridization:Bond angle(s):

Shape: trigonal bipyramidalShared electrons (bonds): 5Lone pairs of electrons: 0Hybridization: sp3dBond angle(s): 90 and 120

Shape:Shared electrons (bonds): 4Lone pairs of electrons: 1Hybridization:Bond angle(s):

Shape: see-sawShared electrons (bonds): 4Lone pairs of electrons: 1Hybridization: sp3dBond angle(s): 90> & 120>

Shape:Shared electrons (bonds): 3Lone pairs of electrons: 2Hybridization:Bond angle(s):

Shape: T-shapedShared electrons (bonds): 4Lone pairs of electrons: 1Hybridization: sp3dBond angle(s): 90> & 120>

Shape:Shared electrons (bonds): 2Lone pairs of electrons: 3Hybridization:Bond angle(s):

Shape: linearShared electrons (bonds): 4 Lone pairs of electrons: 1Hybridization: sp3dBond angle(s): 180

Shape:Shared electrons (bonds): 6Lone pairs of electrons: 0Hybridization:Bond angle(s):

Shape: octahedralShared electrons (bonds): 6Lone pairs of electrons: 0Hybridization: sp3d2Bond angle(s): 90

Shape:Shared electrons (bonds): 5Lone pairs of electrons: 1Hybridization:Bond angle(s):

Shape: square pyramidShared electrons (bonds): 5Lone pairs of electrons: 1Hybridization: sp3d2Bond angle(s): 90>

Shape:Shared electrons (bonds): 4Lone pairs of electrons: 2Hybridization:Bond angle(s):

Shape: square planarShared electrons (bonds): 4Lone pairs of electrons: 2Hybridization: sp3d2Bond angle(s): 90

__ is the combining of two or more orbitals of nearly equal energy within the same atom into orbitals of equal energy

Hybridization

In the case of methane, they call the hybridization sp3, meaning that an __ orbital is combined with three __orbitals to create __ equal hybrid orbitals. These new orbitals have slightly __ energy than the 2s orbital and slightly energy than the 2p orbitals.

spfourmoreless

sp hybridization, in which one _ orbital combines with a single _ orbital. This produces _ hybrid orbitals, while leaving _ normal p orbitals

sp2two

Another hybrid is the sp2, which combines two orbitals from a _ sublevel with one orbital from an _ sublevel. _ p orbital remains unchanged.

psOne

attractions between moleculesoccur within a molecule and include covalent bonds b/w atoms

intermolecular forcesintramolecular forces

Intermolecular attractions:1) attractions b/w opposite charges, strongest kind2) attraction b/w 2 polar molecules that increases with higher polarity3) special type of dipole-dipole b/w H, F, O or N4) occur in all substances, very weak

1) ion-ion2) dipole-dipole3) hydrogen bonding4) london dispersion forces

Put the four types of intermolecular attractions in order from strongest to weakest

ion-ion, hydrogen bonding, dipole-dipole, dispersion

Flame Tests: Red/orange: __, orange: __, orange/yellow: __, magenta: __, yellow/green: __, green/blue: __, violet: __, white: ___

Red/orange: strontium, orange: calcium, orange/yellow: sodium, magenta: lithium, yellow/green: barium, green/blue: copper, violet: potassium, white: magenesium

Single bond: _ sigma bond(s), _ pi bond(s)Double bond: _ sigma bond(s), _ pi bond(s)Triple bond: _ sigma bond(s), _ pi bond(s)

Single bond: 1 sigma bond(s), 0 pi bond(s)Double bond: 1 sigma bond(s), 1 pi bond(s)Triple bond: 1 sigma bond(s), 2 pi bond(s)

Bond order:0: __ 1: __2: __3: __

0: doesn't exist1: single bond2: double bond3: triple bond