chem unit 2 (1.7, 1.8, 2.1-2.7, 3.1)

atomic radius trend across period and why

decreases because more protons so higher Z_eff and nuclear force

atomic radius trend down a group and why

increases because more electron shells

ionic radius trend for cations

smaller than parent atoms because they lose valence shells which causes nuclear force to increase and also causes less electron-electron repulsion

ionic radius trend for anions

larger than parent atom because electrons added to valence shell so theres more electron-electron repulsion and thus larger radius

ionization energy definition

energy required to remove valence electron from gaseous neutral atom in ground state

ionization energy trend across period and why

increases because smaller radius so greater nuclear force which makes it harder to remove electrons

ionization energy trend down a group and why

decreases because nuclear force is smaller (larger radius) and more shells so more shielding; easier to pull off electrons

when are there exceptions to the ionization energy trends

when removing electrons makes the atom more stable

determining amt of valence electrons from successive ionization energies

when there’s a huge jump in IE, you’ve pulled off the last valence electron

electron affinity definition

neutral atom’s liklihood of gaining an electron (but with amount of energy involved); usually neg values

electron affinity trend across period and why

increases because atom gets more stable as you add electrons and thus releases more energy

electron affinity trend down a group and why

decreasing (only for group 1) because larger radius means new electrons are added further from nucleus and thus release more energy when added (doesn’t make it a TON more stable)

what kind of electron affinity values do noble gases have

positive values because adding electrons makes them less stable

electronegativity definition

ability of atom to attract shared electrons

electronegativity trend across period and why

increases because atomic radius decreases while Z_eff increases (nuclear force stronger)

electronegativity down a group and why

decreases because further from nucleus and thus more shielding

what is used to determine polarity

electronegativity

what kind of elements have similar chemical properties

elements with the same valence configuration (same group)

how would you explain periodic trends or anything really?

electron arrangement, proton number, or coulomb’s law

how valence electrons work in ionic bonds

electrons go to element thats the most electronegative; held together by electrostatic force between metal and nonmental

bonds and potential energy graph sign meaning

• positive value indicates gained energy relative to reference state (repulsive force)

• negative value indicates released energy relative to reference state (stable; attractive force)

what does the dip in the PE graph represent

where bonded atoms are most stable and average bond length

how do double/triple bonds affect bond length

makes them shorter bc those bonds are stronger and thus require more energy to be broken

bond length is proportional to

radius because as radius increases, there are more electron shells in the way and the length is bigger

which is more important for determining bond length: bond order or radius

bond order

ionic bonds are between

metals and nonmentals; valence electrons transferred

ionic bonds held together by

electrostatic/coulombic attraction

structure of ionic bonds

crystal lattice

ionic bond properties

• hard/brittle

• conducts electricity when dissolved into ions

• high melting and boiling point

how do you figure out the bond strength of ionic bonds

coulomb’s law

EN difference for ionic bonds

large

covalent bonds are between

2 non-metals; shared valence electrons

covalent bond properties

• non-lustrous

• can be brittle, hard, or soft

• poor conductors

• various colors

polar covalent bonds

• electrons shared unequally

• slight EN difference

nonpolar covalent bonds

• electrons shared equally

• low EN difference (C-H bond considered nonpolar)

metallic bonds are between

2 metals; has sea of electrons

metallic bonds hold together

a crystal lattice

metallic bond properties

• shiny/lustrous

• malleable

• ductile

• conductive (heat and electricity)

metals usually

lose electrons to form cations

why is the ionic structure a crystal lattice

maximize attractive forces and minimize repulsive forces

what does strength of ionic bonding/lattice energy come from

coulomb’s force; lattice energy = energy needed to pull ions infinitely apart

e.g. larger F = stronger attraction and higher melting point

alloy definition

mixture of elements with metallic properties (crystal lattices)

interstitial alloy (general)

small atoms filling spaces btwn larger atoms (usually carbon) — holds them together; resists sliding

interstitial alloy properties

• higher melting point than pure version

• less malleable

• harder

• stronger

substitutional alloy (general)

held together by metallic bonds, free electrons, and positive core; has metal atoms that replace similarly sized metal atoms

substitutional alloy properties

• lower melting point than pure version

• corrosion resistant

sigma bonds

in every bond; head to head overlap of atomic orbitals (handshake bond)

pi bonds

happes w/ double or triple bonds; above and below plane of atom (like North and South points)

intermolecular forces

electrostatic attraction/force between separate molecules

IMFs determine

physical properties like boiling point, melting point, and vapor pressure

london dispersion forces (IMF)

temporary instantaneous dipoles from random electron movement

what do LDFs stick together

neutral atoms and nonpolar molecular

LDFS are present in

all molecules but is the only force in nonpolar molecules

strength of LDFS increases with

more electrons because more polarizability

dipole-dipole forces (IMF)

attraction between permanent positive and permanent negative ends of polar molecules

hydrogen bonding (IMF)

occurs when hydrogen bonds to F/O/N; not actual bonds but more of an interaction (essentially dipole-dipole on steroids)

polarity depends on

3d shape of molecule and electron density

polar bonds depend on

electronegativity

polar molecules depend on

shape (non symmetrical means polar)

lone pairs around a central atoms almost always means

polar molecule

polarizability

“squishiness” of electron cloud"; ease at which electron cloud can be distorted by nearby charges

(e.g. more electrons means harder for nucleus to control electrons and thus stronger temporary dipoles)

dipole-induced force

between polar and nonpolar molecule; force strength increases with dipole magnitude of polar molecule and polarizability of nonpolar molecule

dipole-dipole force

between polar molecules; depends on dipole strength and relative orientation (usually greater interaction in polar molecules because dipole-dipole acts in ADDITION to LDFs)

ion-dipole

betwen ions and polar molecules (stronger than dipole-dipole); usually in water

dipoles point towards the more

electronegative ion (which has a partially negative charge)

to determine molecule polarity

YOU HAVE TO DRAW THE LEWIS STRUCTURE AND USE FORMAL CHARGES

boiling point is related to

relative strength of attractive forces between particles

hydrogen bond donor requirements

H atom covalently bonded to F/O/N

hydrogen bond acceptor requirements

substance has F/O/N with lone pair and is part of polar bond

ion-dipole force is proportional to

magnitude of ion charge