• Shuffle
Toggle On
Toggle Off
• Alphabetize
Toggle On
Toggle Off
• Front First
Toggle On
Toggle Off
• Both Sides
Toggle On
Toggle Off
Toggle On
Toggle Off
Front

### How to study your flashcards.

Right/Left arrow keys: Navigate between flashcards.right arrow keyleft arrow key

Up/Down arrow keys: Flip the card between the front and back.down keyup key

H key: Show hint (3rd side).h key

A key: Read text to speech.a key

Play button

Play button

Progress

1/54

Click to flip

### 54 Cards in this Set

• Front
• Back
 trends 1. size and radius of atoms and ions 2. ionization energy 3. electron affinity 4. electronegativity effective nuclear charge (Zeff) the charge experienced by a particular electron in a many electron atom. calculation of effective nuclear charge Zeff = Z - S where Z= atomic #/nuclear charge and S=avergae # of screening e- (ex) for Na Z=11 (b/c 11 protons) S=10 (b/c 10 e- in between) Z - S = 11 - 10 = +1 Trend (1): size of atoms and ions when atoms interact they: 1. collide and ricochet off (no bond) 2. collide and bond trends for size of atoms 1. down a group, atoms become larger: n increases, so distance between nucleus and outermost e- increases so atomic radius and size increases 2. across a period (from L to R) the size decreases b/c the Zeff increases trends for ions 1. cations are smaller than their parent atom 2. anions are larger than their parent Trend (2): Ionization Energy (I) the minimal energy required to remove an e- from a ground state of a gaseous atom or ion. first ionization energy (I1) amount of energy required to remove an e- from a gaseous neutral atom (ex) Na -- Na+ + e- I1=495 kj/mol (get from table) 2nd ionization energy (I2) energy required to remove second e- from an ion not an atom trends for I: 1. generally increases across a period 2. generally decreases as we go down a group exceptions to ionization trends 1. removal of the 1st and 4th p e-; removal of 1st p e- decreases energy b/c better shielding of nucleus in s orbital b/c nucleus in completely surrounded so for ion (s2p0) gives us lower energy = lower I 2. when 4th e- is added to p and paired an increase in energy b/c they repel each other so the removal of the 4th e- = decrease in energy, lower I. Trend (3): Electron Affinity energy change that occurs from adding an e- to a gaseous atom trends for electron affinity 1. increase from L to R across period 2. does not change much as you go up or down a group greatest affinity = low affinities = halogens (7A)b/c only need one e- to get stable nobel gas configuration = lower energy low affinities for 5A due to adding the e- to half filled p shell Trend (4): Electronegativity measure of the ability of an atom in a bond to attract e- to itself trends for electronegativity 1. increases across of period from L to R 2. increases up a group; fluorine is most electronegative element; Cs is least; C and H have almost the same electronegativity group trends 1. metals 2. nonmetals 3. metalloids trend for metalic character decreases from L to R Metals shiny, lustrous, various colors, sold metals are malleable (hammered into thin sheets) and ductile (pulled into thin wires); good conductors of head and electricity; most metal oxides form ionic solic which forms basic solution; tend to form cations in solution nonmetals not lustrous, many colors; solids are usually brittle; some hard and some soft; poor conductors; nonmetal oxides are usually molecular substance which form basic solution; tend to form anions in solution metalloids share properties of metals and nonmentals to varying degrees. family trends: (1) Alkali metals 1. group 1A 2. soft w/ low melting points 3. low I1 wh/ allows them to make +1 ions easily 4. very reactive (a)combine directly with most nonmetals (b)form metal hydrides, metal sulfides, and metal halides (c) react vigorously w/ water family trends: (2) Aldaline Earth Metals 1. group 2A 2. low I1 but not low as 1A's; less reactive 3. the heavier the metal the more reactive examples: 1. Be: no reaction with H2O liquid or with H2O gas 2. Mg: no reaction with H2O liquid but will react low with H2O gas 3. Ca, Sr, Ba: slow reaction with H2O liquid and reaction with H2O gas family trends: (3) Hydrogen 1. normally seen as H2 gas: under extreme pressure can be made metallic 2. combustable 3. prefers to share e- in cavalent bonds in molecular compounds 4. reacts with metals and nonmetals (a)nonmetals = exothermic (b)metals = forms solid metal hydride (c) will lose e- to form H+ seen mostly in sol family trends: (4)Oxygen 1.nonmetallic to metallic as move down group 2. O2 is typical form (other 6A's are solid) 3. ozone (O3) 4. forms O2- ions 5. forms O2 2- peroxide ions and O 2- superoxide ions sulfur 1. most stable = S8 solid family trends: (5) Halogens 1. group 7A 2. nonmetals, diatomic 3. easily gain e- to form -1 ions 4.F2 is most reactive and strongest oxidizer 5. Cl2 is very reactive but less than F2 6. react with most metals to form ionic halides family trends: (6)Nobel Gases 1. group 8A 2. nonmetal 3. highly unreactive; don't readily lose e-; have octet e- configuraton electron configuration for ions 1.cation:remove e- from elemental configuration: remove e- from orbital with largest n. 2.anions: add e- to elemental configuration with lowest available n 3 types of bonding 1. transfer e- form one atom to another: Ionic Bonding 2. sharing e- b/t atoms: Covalent Bonding 3. delocalization of e- over whole solid: Metallic Bonding Lewis symbols way of depicting valence e- for bonding; disperse e- around symbol using hund's rule; for representative elements Bonding Basics 1. octet rule: reactivity and bonding generally work toward an octet or nobel gas configuration 2.nobel gases a full outer shell (8e-) except for He which has 2 e- in full shell 3. other atoms tend to gain or lose e- or share e- to reach this stable full outer shell conf. ionic bonding energy released in reaction forming stable product; bond formed by attraction between positive and negative charges lattice energy use to break up; solid network or crystal lattice formed by ions; lattice energies = large positive #'s Eel = KQ1Q2/d where: k = 8.99 X 10^9 jm/c^2 Q1,Q2 = charge of ions d = distance (m) why do ions form as they do? Na=[Ne]3s^1 Na+=[Ne] stable nobel conf. to remove a 2nd e- comes from inner core; requires more energy than lattice energy provides nonmetals... gain of e- is either exothermic or slightly endothermic if added to a valence shell. example Cl=[Ne]3s^23p^5 Cl=[Ne]3s^23p^6 to add 2nd e- it would have to be added to higher energy shell and energy input would exceed lattice energy covalent bonds sharing e-, usually b/t nonmetals; most substances are covalent; nuclei repel each other; e- repel each other, but nuclei attract the e- from other atom; when atoms come close enough (bonding distance) then the opposing nuclei attracts opposing e- and e- are shared polyatomic ions 2 or more atoms bound together primarily by covalent interactions to form stable group w/ + and - charges covalent bonds and polarity 1. w/ covalent bonds if e- shared equally: non-polar covalent bond 2. covalent bond where e- drawn or attracted to one atom more than another: polar covalent bonds electronegativity scale 1. < 0.5: nonpolar covalent 2. b/t >/= 0.5 & < 2.0: polar covalent 3. >/= 2.0: ionic bond polarity 1. electronegativity used to measure 2. points to more electronegative atom dipole moment measure of polariy of a molecule u = Qr where r = bonding radius formal charges help decide b/t reasonable structures calculating formal charges FC=[#valence e- for atom] - [# nonbonding e- on atom] - [1/2 #bonding e- on atom] reasonable structures they are stable as either and interchange (or resonate) b/t the structures example: ozone exceptions to octet rule 1. molecules with odd # of total valence e- 2. molecules which central atom can exist stably w/ less than an octet 3. molecules where central atom can accomidate more that an octet bond energy strength of bonds determined by the energy necessary to break the bond bond enthalpy always requires energy to break bond; always positive bond length the distance b/t the nuclei of bonding atoms bond order a measure of the # of bonding e- pairs b/t atoms 1.single bond= bond order 1 2.double bond= bond order 2 3. triple bond= bond order 3