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Chapter 8 Chemistry AP

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Section 1

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Covalent Bond (Definition)

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Chemistry Science

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Mar 1, 2020

Cards (30)

Section 1

(30 cards)

Covalent Bond (Definition)

Front

sharing of electrons

Back

Use formal charge to choose correct LS

Front

1) Correct formal charges, and closest to zero 2) Negative charges are on the most electronegative atoms (EN increases up and right on PT)

Back

Dipole moment (in polar covalent bond)

Front

- When two atoms share electrons unequally, a dipole moment, µ, results µ = Qr

Back

Lewis structures (LS)

Front

- Shows how valence electrons are arranged - Dots are electrons, lines are bonds - Atoms follow octet rule (EXCEPT Hydrogen, needs 2)

Back

Bond Energy

Front

Energy required to break a bond

Back

Isoelectronic Ions

Front

These ions have the same number of electrons (same electron configuration. ex) Al+3, Mg+2, Na+1, Ne, F-1, O-2, N-3 (all have 10 electrons) The charges add/subtract electrons to a point where they are all the same. For example, O does not have the same # electrons as Ne, however O-2 does because it has two more electrons than O, just like Ne.

Back

Bond length

Front

As the number of bonds between two atoms increases, the bond length decreases

Back

How to calculate ΔH with lattice energy

Front

Add up all energies in reaction (sublimation of elements, ionization, dissociation, formation of ions, etc) then add the ionization energy (always negative) which represents the formation of the product (answer for ΔH should be negative)

Back

Ions

Front

- Metals lose electrons to form cations - Nonmetals share electrons in covalent bonds (two nonmetals) or gain electrons to form anions in ionic bond

Back

Electronegativity difference

Front

As difference in EN increases, bond type changes from non polar covalent (EN=0-.4) to polar covalent (EN=0.4-2.0) to ionic (EN>2).

Back

Lattice energy trends

Front

- Lattice energy increases with the charge on the ions - Increases with the decreasing size of the ion

Back

Nonpolar covalent bonds

Front

- Electrons are shared equally - EN difference: 0 - 0.4

Back

Polar covalent bonds

Front

- Electrons are NOT shared equally - EN difference: 0.4 - 2.0 - The greater the difference in EN, the more polar the bond - One end is slightly positive, the other slightly negative (ends indicated with 𝛿+ and 𝛿-, respectively) - The side with 𝛿- has a higher EN

Back

Metallic Bond (Definition)

Front

Metal atoms bonded to several other atoms

Back

Enthalpies of reactions

Front

ΔH(rxn) = (ΔH bonds broken) - (ΔH bonds formed)

Back

Lattice energy

Front

- Energy required to completely separate a mole of solid ionic compound into its gaseous ions (this is required because there is an electrostatic attractions between cation and anion) - Calculated with coulomb's law E(el) = (kQ1Q2) / (d)

Back

Size of isoelectronic ions

Front

Use common sense. Positive ions have more protons so they are smaller, negative charged ions have more electrons so they are larger (b/c force b/w electrons increases, pushing out)

Back

Size of ions

Front

- Cations are smaller than original atom - Anions are larger - Ion size increases down a group - From left to right they get smaller (because nuclear charge increases), then suddenly larger (once it changes from cations to anions)

Back

Formal charge (definition)

Front

Charge atom would have if all the atoms in the molecule had the same EN

Back

Covalent bonding

Front

- Atoms share electrons - Attractions: between electrons and nuclei - Repulsions: between electrons, and between nuclei

Back

Electronegativity (EN)

Front

Ability of atoms IN A MOLECULE to attract electrons to themselves.

Back

How to Draw LS

Front

1) Find sum of valence electrons from atoms in molecule - If there are charges: Anion: add one electron for each negative charge, cation: subtract one electron 2) Write symbols for atoms and connect with single bonds (lines) - Central atom is the least electronegative element that isn't hydrogen 3) Fill octets with pairs of dots 4) Place remaining electrons around central atom (even if it has more than an octet) 5) If you run out of electrons before the central atom has an octet, try using multiple bonds.

Back

Ionic Bond (Definition)

Front

electrostatic attraction between ions

Back

Octet rule

Front

When forming compounds, atoms like to be surrounded with 8 valence electrons

Back

Resonance structures

Front

- One LS does not adequately describe molecule - Must draw two LS - The electrons can move among the elements - They are delocalized - Circle in hexagon signifies delocalized ions

Back

How to assign formal charge

Front

FC = (Valence electrons) - (Lone pairs for that atoms + 1/2 bonding electrons)

Back

Bond enthalpy

Front

energy required to break a bond (number is given)

Back

Ionic Bonding

Front

An atom with low IE rects with an atom with high EA (usually a metal and nonmetal). - electron moves from metal to nonmetal - opposite charges hold the atoms together (cation and anion)

Back

Exceptions to Octet rule

Front

1) Molecules or ions with odd number of electrons - single electron dot (not a pair) would be placed on one of the atoms 2) Molecules or ions with less than an octet - Charges would be messed up with 8 valence around each element, so put less electrons around center element. Ex) BF3 has 6 valence electrons around Boron, not 8. 3) Molecules or ions with more than 8 valence electrons (expended octet) - Octet expands with third row or below only - Usually center element would have more bonds

Back

Ionic compounds

Front

- Ions align themselves to maximize attraction between opposite charges and minimize repulsion between like ions - Stabilize ions that would be unstable as a gas

Back