U1: trends + bonding

Question 1

prefixes

Answer 1

mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca

Question 2

roman numerals?

Answer 2

I, II, III, IV, V, VI, VII, VIII, IX, X

Question 3

annhydruous

Answer 3

hydrate removed

Question 4

boron, carbon, & silicon charge?

Answer 4

boron: 3+
carbon: 4+
silicon: 4+

Question 5

which periods (rows) can have 8-18 valence electrons?

Answer 5

period 3 and below

Question 6

EN values mean which bonds?

Answer 6

0-0.5 is non-polar covalent
0.5-1.7 is polar covalent
1.7-3.3 is ionic

Question 7

melting and boiling points for each intramolecular bond?

Answer 7

ionic: high
polar-covalent: medium
non-polar covalent: low

Question 8

water solubility for each intramolecular bond?

Answer 8

ionic: water-soluble
polar-covalent: water-soluble
non-polar covalent: non-water-soluble

Question 9

electrical conductivity for each intramolecular bond?

Answer 9

ionic: yes, electrolyte
polar-covalent: no, non-electrolyte
non-polar covalent: no, non-electrolyte

Question 10

electrolyte

Answer 10

substances that conduct electricity in the molten state or in solution

Question 11

when are imfs broken/formed?

Answer 11

broken during boiling/melting, formed during freezing/condensing

Question 12

which properties do imfs impact?

Answer 12

physical. i.e. changes of state, solubility

Question 13

which states are imfs strong in?

Answer 13

liquids + solids. almost non-existent in gases except extreme conditions. (tht’s why there r gas laws)

Question 14

relationship between imf and state at room temp?

Answer 14

strong imf = solid at room temp
weak imf = gas at room temp

Question 15

order from strong-weak of imfs

Answer 15

1. ionic/covalent bonds
2. ion-dipole (ion+dipole)
3. hydrogen bonds (H w/ NOF)
4. dipole-dipole (polar molecules)
5. london dispersion (non-polar, but exist in all molecules)

Question 16

polarizability + what impacts it

Answer 16

“ease with which an electron cloud can be deformed.” larger (more electrons) = more polarizable

Question 17

PEN rule

Answer 17

protons = atomic number
electrons = atomic number - charge
neutrons = atomic mass - atomic number

Question 18

atomic mass unit

Answer 18

• calculated w/ magnetic field
• approx. mass of 1 proton
• 1/12 mass of C-12 atom

Question 19

isotope

Answer 19

• same atom have diff. amnts of neutrons/protons
• each occur in abundance, e.g. 12% means 12% will be that isotope in a sample

Question 20

radioisotope

Answer 20

• has unstable nucleus
• has radioactive decay

Question 21

radioactive decay

Answer 21

• process by which an unstable nucleus rearranges itself
• nucleus can throw out protons/neutrons that are imbalanced

Question 22

electron configuation + wht it’s based off of

Answer 22

• an atoms arrangement of electrons around the nucleus to have the lowest possible energy bc thats the most stable state
• stability based on atomic radii, ionization energy, electron affinities, and electronegativity

Question 23

atomic radii

Answer 23

• size of atom
• measured in picometres (pm) or angstroms (A w/ degree) from centre of nucleus to outermost electron

Question 24

wht’s atomic radii influenced by?

Answer 24

1. energy level: high energy levels are further from nucleus
2. charge on nucleus: more protons=more pull on electrons
3. shielding effect

Question 25

shielding effect

Answer 25

electrons in inner shells (kernel electrons) shield the valence electrons from the positive charge of nucleus, so the valence electrons can move further out = bigger atom. +like charges repel, so kernel electrons repel the valence

Question 26

ions

Answer 26

charged atoms, happen when there's more or less electrons than protons (imbalance)

Question 27

cation

Answer 27

- positively charged ions formed as a result of loss of electron(s) - metals usually form cations -smaller than neutral atom bc. less electrons mean the positive nucleus is more potent

Question 28

anion

Answer 28

- negatively charged ions formed as a result of gain of electron(s) - non-metals usually form anions -larger than neutral atom bc. extra electrons repel each other

Question 29

ionization energy | (wht does a low IE mean)

Answer 29

- amount of energy required to completely remove an electron from a gaseous atom - low IE = easy to take electron = more reactive

Question 30

what impacts ionization energy | (wht causes big/small IE)

Answer 30

- bigger IE w/ bigger charge of nucleus - smaller IE if valence electrons are far away from nucleus - exception: filled and half-filled (4 & 8) valence shells have high IE - shielding effect

Question 31

electron affinity

Answer 31

- energy change associated with adding an electron to a gaseous atom: typically negative due to release of energy - easiest to add to halogens

Question 32

wht does a negative electron affinity cause

Answer 32

atom becomes more stable and thus releases a great deal of energy when it gains an electron - decrease in potential energy

Question 33

wht does a positive electron affinity cause

Answer 33

atom becomes less stable and thus gains energy when it gains an electron - increase in potential energy

Question 34

which group has the highest electron affinity?

Answer 34

halogens: release the most energy when they gain an electron, meaning they become very stable

Question 35

electronegativity

Answer 35

tendency for an atom to attract electrons to itself when it is chemically combined with another element

Question 36

relationship between electronegativity and electron affinity?

Answer 36

atoms with large negative electron affinity have large electronegativity: large electronegativity means it pulls the electron toward it

Question 37

trends of each periodic trend?

Answer 37

electron affinity, ionization energy, electronegativity: increase right and up atomic radius: increase left and down