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Chemistry 1 > Test 6- Bonding > Flashcards

Test 6- Bonding Flashcards

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1
Q

Why do bonds form?

A

Energy! Atoms strive to be at lowest energy possible

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2
Q

Diagram that shows relationship between potential energy level and distance between atoms

A

Be able to draw

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3
Q

Chemical bond

A

Mutual electrical attraction between nuclei and valence (outer) electrons of different atoms. This binds the atoms together. All about valence electrons!

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4
Q

Model to show chemical bonding

A

Lewis dot diagram shows elemental symbol and bonding VE

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5
Q

Lewis Dot diagram basics

A 
One dot per VE
--36
4???1
7???2
--58
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6
Q

Stable electron configuration

A
  • Same as nobel gas

- Elements lose or gain electrons to have the same number of electrons as a nobel gas

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7
Q

Octet rule

A
  • Elements strive for 8e- in outer energy level.
  • some only need 2 e- to be stable like He
  • Elements in part A of periodic table
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8
Q

Exceptions to octet rule

A
  • H only needs 2 e-
  • B needs 6 e-
  • Be needs 4 e-
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9
Q

Electronegativity

A

Relative scale 0-4, the attraction an atom has for a pair of e- in a bond

  • Increases up and right
  • Electrons close to each other on periodic table share electron to meet the octet rule (covalent)
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10
Q

Share electrons

A
  • Similar electronegativity
  • same area of the periodic table
  • share e- to fufil octet rule
  • covalent bond
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11
Q

Transfer electrons

A
  • Different electronegativity
  • far apart on the periodic table
  • Ionic bond
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12
Q

Nonpolar covalent bond

A
  • Perfect sharing
  • Same attraction
  • same electronegativity
  • Electronegativity difference Range: 0-.4
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13
Q

Polar covalent bond

A
  • Unequal sharing

- Range: .4<_1.9

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14
Q

Ionic bond

A
  • Not sharing
  • Complete transfer
  • Range: above 1.9
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15
Q

Drawing Ionic bonds

A
  1. Draw lewis dot for both elements
  2. transfer to make 8
  3. Write both elements with brackets
  4. Indicate the charge
    - Might have to draw more than one to make 8
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16
Q

Characteristics of Ionic compounds

A
  • Simplest piece is a Formula unit
  • Crystal lattice AKA ionic latice
  • Always solid at room temperature
  • Soluble in water
  • Conducts electricity
  • High melting point
17
Q

Formula Unit

A

Simplest ration of the ions

18
Q

Crystal lattice AKA ionic latice

A 
- Cluster of Na+ and Cl-
Na+Cl-Na+Cl-
Cl-Na+Cl-Na+
Na+Cl-Na+Cl-
- Not smooth
19
Q

Ionic compounds solids at room temperature

A

So strongly attracted that it takes a lot of energy to separate

20
Q

Ionic compounds soluble in water

A

Water has a great attraction to the compound and the high energy breaks them apart. (Brad Pitt and Angelina Jolie)
- NaCl(s) –> NaCl(aq)

21
Q

Ionic compounds conducting electricity

A
  • Solution of ions and liquid ionic compounds
  • has to be s solution or melted/liquid
  • Triangle of battery, light, and beaker with solution with wires going into the solution
22
Q

Nomenclature

A
  • Naming/writing formulas

- IUPAC- International Union of Pure and Applied Chemists

23
Q

Naming Ionic compounds

A
  • Binary
  • Name metal first
  • When writing a formula, the net charge for all compounds is zero
  • Subscript indicates number of atoms/ions of the element it follows
  • Criss cross- use charge as subscript of other element
24
Q

Binary

A
  • Ionic
  • 2 elements
  • Metal and nonmetal
  • End in ide
  • name nonmetal and change ending to ide
25
Q

Naming transitional metals

A

Use roman numerals to indicate charge of metal ion

26
Q

Tertiary Ionic compounds

A
  • More than 2 elements
  • cation(+)/anion(-)
  • Don’t end in ide
  • All contain polyatomic ions
27
Q

Naming representative metals

A
  • more than 2 elements
  • Name metal and name polyatomic ion
  • Subscript doesn’t matter
  • Not all compounds that end in Ide are not binary but all binary compounds end in ide
  • If you have more than one polyatomic ion, it must be in parenthesis
28
Q

Transitional metals with polyatomic ions

A

Use roman numerals to indicate charge

29
Q

Covalent bonds

A
  • Occur between 2 nonmetals (in general)
  • 2 types of covalent bonds based on electronegativity difference
  • 0-.4= nonpolar covalent bond, equal sharing
  • 0<_1.9= polar covalent bond unequal sharing
  • Exist as individual molecules
  • Molecule=smallest piece of covalently bonded particle
30
Q

Smallest piece of covalent vs. Ionic

A
  • Covalent= molecule

- Ionic= Formula unit

31
Q

Representing e- sharing covalent bond with a lewis dot diagram

A
  • Still all atoms strive for 8e-
  • No transfer
  • No charges or brackets
  • Dots and circles
  • Covalent bond is the overlap
  • e- always shared in pairs
  • Single if one pair is shared
  • Double if two pairs are shared
  • Triple if three
32
Q

Line structure

A

Each line represents a pair of electrons

33
Q

Lewis Dot for polyatomic ions

A
  • have circles (2 nonmetals) and brackets (charge)
  • Lower electronegativity goes in the middle
  • Add/subtract electrons according to charge
34
Q

Coordinate covalent/Dative bonds

A

Both shared electrons are from one element

35
Q

Lewis Dot and Line are both

A

structural diagrams

Chemistry 1 (17 decks)

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