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Chem 11 > Periodic Table/trends > Flashcards

Periodic Table/trends Flashcards

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0
Q

Henry Moseley

A

Was 1st to arrange elements in a table format by atomic NUMBER

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1
Q

Dimitri Mendeleev

A

Produced 1st periodic table

  • arranged elements according to atomic MASSES
  • his table had empty spaces for elements yet to be discovered
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2
Q

Periodic law

A

When elements are arranged by atomic number, similar properties will exist in every 8 or so elements

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3
Q

Rows are called

A

Periods

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4
Q

Columns are called

A

Families or groups

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5
Q

Alkali metals

A
  • 1st column/group
  • reactive with water and air
  • generally a 1+ charge
  • soft metals (can cut with a butter knife)
  • forms basic solutions
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6
Q

Alkaline earth metals

A
  • 2nd column/group
  • not as reactive as Alkali metals
  • generally 2+ charge
  • also somewhat soft metals
  • forms basic solutions
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7
Q

Halogens

A
  • 2nd last column/group
  • very reactive non-metals
  • generally a -1 charge
  • F (g) Cl (g) Br (l) I (s) At (s) diff. phases
  • forms acidic solution
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8
Q

Noble gases

A
  • generally non-reactive (inert) gases
    (some exceptions)
  • full outer shell
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9
Q

Transition metals

A
  • d block
  • metals in this group have MANY different charges
  • results in colored solution
  • tend to form metal complexes w/ certain ligands (attachments)
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10
Q

Metallic properties

A

Shiny, luste, malleable, ductile, conductive ect.

|———–non-metal
| I
|. I
Metal——————

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11
Q

Atomic size (radii)

A
  • going down a column INCREASES atomic size. (orbitals get larger)
  • going across from Left to Right, atomic size DECREASES
    - if # protons increases it leads to a greater nuclear charge; can attract electrons closer in
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12
Q

Ion size (compared to neutral atoms)

A
  • anions (-) will be LARGER than their neutral atom
    - more e. w/ same # protons; smaller pull per e.
  • cations (+) will be SMALLER than their neutral atom
    - fewer e. w/ same # protons; greater pull per e.
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13
Q

Ionization energy (IE)

A
  • energy required to remove an outermost e. from a atom
    - 1st IE -> energy required to remove first outermost e.
    • 2nd IE -> energy required to remove second outermost e.

DECREASES going down
- because e. are farther away from nucleus
INCREASES left to right
- because of greater nuclear charge (larger pull)

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14
Q

Electronegativity

A
  • the ability of a nucleus to attract electrons towards its self WITHIN A BOND

DECREASES going down

   - outermost e. is further away
         - > Shielding effect

INCREASES going left to right
- more protons; greater nuclear pull

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15
Q

Shielding effect

A

Inner e. will “shield” away some of the nuclear pull towards outer e.

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16
Q

Remember:

A

SIZE OVERPOWERS, # protons in the nucleus
- look at size first

Look up/down before left/right

These are trends - many exceptions exist

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17
Q

Lanthanides/Actinides

A

Radioactive

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18
Q

Radioactive elements

A

Lanthanides

Actinides

19
Q

Types of chemical bonding

A

Ionic, covalent, polar covalent

20
Q

Ionic bonds

A

Results from the electrostatic attraction between cations and anions

21
Q

Mechanism of ionic bonds

A

1) an atom (usually a metal) transfers electron(s) to another atom (usually a non metal)
2) the atom that lost e. becomes + charged (cation) while the atom that gained e. becomes - charged (anion)
3) the cation will be attracted to the anion through electrostatic attraction (opposite charges attract)

22
Q

Covalent and polar covalent similarities

A

Arise from a sharing of e.

23
Q

Covalent and polar covalent defferences

A

Pure covalent

  • equal sharing of the bonding pair of e.
  • similar Electronegativity values

Polar covalent

  • unequal sharing of the bonding pair e.
  • sig. Different Electronegativity values between atoms
24
Q

Polar covalent

A
  • unequal sharing of the bonding pair e.

- sig. Different Electronegativity values between atoms

25
Q

Pure covalent

A

Pure covalent

  • equal sharing of the bonding pair of e.
  • similar Electronegativity values
26
Q

Its ionic if the Electronegativity value is

A

Greater than 1.7

27
Q

It’s polar covalent if the Electronegativity value is

A

Between 0.5 and 1.7

28
Q

It’s pure covalent if the Electronegativity value is

A

Less than 0.5

29
Q

In a covalent compound which atoms MUST Have 8 e around them

A

C, N, O, F

30
Q

Which atom only needs 6 e.

A

B

31
Q

VSEPR stands for

A

Valence Shell Electron Pair Repulsion

32
Q

Polarity of molecules

A

If these polar bonds in a molecule do NOT cancel each; then the molecule is considered polar (net dipole)

33
Q

Noun of polar

A

Net dipole

34
Q

Solubility - like dissolves like

A

Non-polar solutes dissolve in non polar solvents

polar solutes dissolve in polar solvents

35
Q

What VSEPR shapes are always polar (due to the location of the lone pairs)

A 
Bent/angular
Trigonal pyramid 
Seesaw
T shape
Square pyramid
36
Q

Intermolecular forces (IHF)’s

A

Attractions BETWEEN molecules

37
Q

Intramolecular forces

A

Ionic, covalent and polar covalent

within

38
Q

6 types of Intermolecular forces (IHF’s)

Weakest to strongest

A
  1. London Dispersion forces/Van der Waals forces
  2. Dipole - Dipole attractions
  3. Hydrogen Bonding
  4. Metallic bonding
  5. Ionic Lattice structures
  6. Covalent network solids
39
Q

London dispersion forces/ Van der Waals Forces

A

Weakest

  • attractions that arise form TEMPORARY dipoles in molecules
  • exist for all chemical species
40
Q

Dipole - Dipole attractions

A

2nd weakest

  • exist between POLAR MOLECULES (ie. have PERMANENT dipole)
  • therefor it’s a bit stronger than LD forces
41
Q

Hydrogen bonding

A

3rd weakest

  • attraction between H of one molecule and a lone pair of a
    N, O, F
    Atom of a second molecule
42
Q

Metallic Bonding

A

4th weakest

Metal atoms are attracted to each other and can pass e. around (explains why metals are good conductors)

43
Q

Ionic lattice structure

A
  • recall: ionic compounds exist as crystal lattices

- atoms in this structure are STRONGLY attracted to each other

44
Q

Covalent network solids

A
  • like ionic lattices but w/ covalent bonds

Ex. Carbon (graphite), Carbon (diamond), SiO2 (quartz)

Sometimes: metal carbides and silicides

45
Q

Connection to KMT/ boiling melting point

A
  • energy required to melt/boil a substance is used to counteract the IMF’s present
  • substances with STRONG IMF will have HIGH melting/boiling point

Chem 11 (12 decks)

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