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Chem 100 Class > Unit 4 > Flashcards

Unit 4 Flashcards

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1
Q

Main group metal cations are named by identifying the metal, followed by the word ion.

A

• Recognizing polyatomic ions in a chemical formula helps to name them correctly.

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2
Q

Naming cations

A

Metals become (+) ions = cations; keep name of element + the word “ion”
• Ex. Na+ , sodium ion

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3
Q

Naming anions

A

Non-metals become (-) ions = anions; keep 1st syllable of element name, drop ending and add suffix –ide
• Ex. Cl-, chloride ion

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4
Q

Some metal and nonmetal ion names

A
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5
Q

[metals with variable charge]
Transition metals form two or more positive ions (cations). – except for

A

Zn2+, Cd2+, and Ag+

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6
Q

[metals with variable charge]
Roman numeral

A

-A Roman numeral equal to the ion charge is placed in parentheses immediately after the metal name.
-Cu2+ copper(II)
-Cu+ copper(I)
-Fe2+ iron(II)
-Fe3+ iron(III)
-Pb2+ lead(II)
-Pb4+ lead(IV)
-Cr2+ chromium(II)
-Cr3+ chromium(III)

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7
Q

[metals with variable charge]
Names of some transition metal cations

A
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8
Q

[Polyatomic ions]

A

a group of atoms with an overall net charge (+ or −).
• Atoms covalently bound—share electrons
• often consist of nonmetals such as P,S,C,N,O and H.
• usually have a negative charge (1−, 2−, or 3−)
• Sum of all e- > sum of all of protons in the combined
atoms. (# e- > # protons)
• except for NH4+, ammonium, which has a (1+) charge
• (#protons > # e-)

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9
Q

Names of Polyatomic ions

A

Names of most common polyatomic ions end in ate.
SO42− (sulfate)
PO43− (phosphate)
NO3− (nitrate)
a. When a related ion has one less oxygen, its name ends in ite.
SO32− (sulfite)
PO33− (phosphite)
NO2− (nitrite)
b. Exceptions to these rules are the following:
CN− (cyanide)
OH− (hydroxide)

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10
Q

Some common Polyatomic ions

A
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11
Q

Naming Ionic compounds

A

Ionic compounds are named by citing first the cation and then the anion, with a space between words.

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12
Q

Flow chart, naming IONIC compounds

A
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13
Q

[naming ionic compounds]
Type I ionic compounds

A

Cations of main group elements.
• charges do not vary.
• Do not specify the charge on the cation.
NaCl is sodium chloride.
MgCO3 is magnesium carbonate.

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14
Q

[naming ionic compounds]
Type II ionic compounds

A

Cations of transition metals
• exhibit more than one charge.
• Specify the charge on the cation in these compounds
**
FeCl2 is iron(II) chloride or ferrous chloride.
FeCl3 is iron(III) chloride or ferric chloride.

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15
Q

Naming covalent compounds

A

-2 non-metals
-first nonmetal
Use a prefix to match a subscript before the element name
-second nonmetal
Use a prefix to match a subscript before the element name and end with ide

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16
Q

Prefixes used to name molecular compounds

A

1 mono
2 di
3 tri
4 Tetra
5 penta
6 hexa
7 hepta
8 Octa
9 nona
10 deca

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17
Q

Naming molecular covalent compounds based on formula

A

• 1st nonmetal – its element name.
(Least electronegative element)
• 2nd nonmetal –first syllable of
the name followed by ide.
• Add prefix to represent subscripts to the names of atoms
• indicates two or more atoms of an element
• doesn’t matter if it is the 1st or 2nd atom in formula

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18
Q

When sodium or other metal reacts with chlorine or any other halogen or nonmetal

A

the metal transfers an electron from its valence shell to the valence shell of the halogen or nonmetal.

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19
Q

Transition metals

A

Form cations
-the metal transfers an electron from its valence shell to the valence shell of the halogen or nonmetal.

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20
Q

Ionization energy

A

• energy absorbed to remove one electron from a single
atom in the gaseous state.
– low values lose electron easily to form cations.

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21
Q

Electron affinity

A

• energy released on adding an electron to a single atom
in the gaseous state.
– large values—>gain electrons easily to form anions

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22
Q

Ion formation

A

• Halogens gain electrons most easily.
• Alkali metals lose electrons most easily.
• Elements near the middle of the periodic table do not form ions easily.
• Noble gases do not form ions.

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23
Q

Rules for writing formulas for ionic compounds:

A
  1. Determine the numbers and kinds of ions in the compound
  2. List the cation first and the anion second.
  3. Do not write the charges of the ions.
  4. Use parentheses around a polyatomic ion formula if it has a subscript.
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24
Q

Writing formulas with Polyatomic ions

A
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25
Q

Properties of ionic compounds

A

• consist of positive and negative ions
• ionic bonds– electrostatic attractions between (+) and (-) ions
• Ion-transfer reactions of metals and nonmetals form products unlike either element.
• holds ions together as crystals in an ionic solid
• high melting and boiling points
• solids at room temperature.

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26
Q

Ionic solids

A

• Crystal lattice
– Ions in each compound held rigidly in place by attraction to their neighbors.
• Ions settle into a pattern that efficiently fills space and maximizes ionic bonding.
• Ionic solids shatter if struck sharply.

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27
Q

Covalent bond

A

-A bond formed by sharing electrons between atoms
-Molecular compounds form when
• atoms of two or more nonmetals share electrons and form
a covalent bond.
• valence electrons are shared between nonmetal atoms to
achieve noble gas electron configuration.
A molecule forms when two or more atoms share electrons.
– A molecule is a type of compound

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28
Q

Diatomic molecules

A

• e- are shared equally
• There are seven diatomic elements: H2, N2, O2, F2, Cl2, Br, and I2.
– Covalent bonding in hydrogen (H2):
– Spherical 1s orbitals overlap
» 1 e- from each atom
– providing 1s2 configuration of He for
each H atom
– H-H, H:H, and H2 all represent a hydrogen molecule.
– Covalent bonding in diatomic molecules such as Chlorine (Cl2)
– overlap of p orbitals.
» 1 e- from each atom
– For Cl: 3p orbitals overlap
» each Cl atom has configuration
1s2 2s2 2p6 3s2 3p6 of Ar
– Cl-Cl, Cl:Cl, and Cl2 all represent a chlorine molecule.

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29
Q

Two types of covalent bonds:

A

Typical covalent bond- each atom donates an electron to form the bond
Coordinate covalent bond– both electrons are donated by the same atom to form the bond
– Explains the bonding…
in polyatomic ions (ex. NH4+ and H3O+)
between H+ (protons) and negatively charged polyatomic ions (bases) in the formation of acids

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30
Q

Covalent bonds:

A

• Between atoms of nonmetals
• shared valence electrons.
• Form molecular compounds or molecules
• Non-electrolytes

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31
Q

Ionic:

A

• Mainly between metal and nonmetal atoms
• valenceelectronstransfer from the metal atom to the nonmetal atom.
• Form ionic compounds or
salts
• Most salts are electrolytes
• Acids and bases are made up of ions

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32
Q

Covalent bonds and the periodic table

A

of e- a nonmetal atom shares usually equals # …

• of e- it needs to achieve a stable e- configuration
• of covalent bonds it forms

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33
Q

Exceptions to the octet rule

A

Boron
– has only three electrons to share, and forms compounds in which it has only three covalent bonds and six valence electrons.

Elements in the third row and below
– have vacant d orbitals that can be used for bonding, allowing them to have an expanded octet.

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34
Q

Multiple covalent bonds

A

Carbon, nitrogen, and oxygen are the elements most
often present in multiple bonds.
• C and N can form double and triple bonds. • O can forms double bonds.
The only way for the atoms in CO2 and N2 to have outer-shell electron octets is by sharing more than two electrons.

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35
Q

Characteristics of molecular compounds

A

-Molecules are neutral, so there is no strong electrostatic attraction between molecules.

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36
Q

Ionic vs. Molecular Formulas

A

• molecular formula: representation of # and kinds of atoms that are
combined in one molecule.
• ionic formula: represents a ratio of ions.

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37
Q

Drawing Lewis Structures Of Molecules And Polyatomic Ions

A

Things to remember while drawing Lewis structures
• Using the chemical formulas to write the atoms in a skeletal structure of the compound
– Don’t draw dot structures for each atom and connect them – The central atom will be:
• The least electronegative atom OR atom that can form the most bonds
– Terminal positions will be occupied by: • Hydrogen–always
• Halogens– usually
– except when more electronegative atoms present in compound
– More than one carbon will form chains of carbon-carbon covalent bonds

38
Q

VESPR Theory

A

Valence Shell Electron-Pair Repulsion Theory (VSEPR)
• describes the orientation of electron groups/clouds around the
central atom. • Statesthat…
• electron groups/clouds are arranged as far apart as possible around the central atom to minimize repulsion between their negative charges—>molecular geometry
• the specific shape of a molecule is determined by the number of atoms attached to the central atom—>molecular shape

39
Q

Applying the VSEPR model

A

STEP 1: Draw a Lewis structure of the molecule, and identify the atom whose geometry is of interest.
• This is usually the central atom.
STEP 2: Count the number of electron charge clouds surrounding the atom of interest.
• This is the total number of lone pairs plus shared e- pairs.
STEP 3: Predict molecular geometry or shape by assuming that the charge clouds orient in space so that they are as far away from one another as possible.

40
Q

Molecular geometry and shapes

A
41
Q

Linear angle

A

180 degrees

42
Q

Trigonal planar

A

120 degrees

43
Q

Trigonal planar, bent
2 bonds 1 lone pair

A

> 120 degrees

44
Q

Tetrahedron, tetrahedral

A

109 degrees

45
Q

Tetrahedron, Trigonal pyramidal
3 bonds 1 lone pair

A

> 109 degrees

46
Q

Tetrahedron, bent
2 bonds
2 lone pairs

A

> 109

47
Q

Polarity of Bonds

A

Nonpolar bond – electrons are roughly shared equally
– Ex. In the H2 molecule (pure non-covalent bond)
or CH4 molecule
• Polar bond – electrons are shared unequally
– One nucleus holds to e- pair closer than the other
– Ex. In the HCl molecule

48
Q

Electronegativity

A

The electronegativity of an atom is its ability to attract the shared electrons in a covalent bond. It…
• from left to right going across a period on the periodic table.
• from bottom to top of the periodic table.
• High for the non metals (F is the highest)
• Low for the metals.

49
Q

Polar covalent bonds and electronegativity

A

The electronegativity of an atom is its ability to attract the shared electrons in a covalent bond. It…
• from left to right going across a period on the periodic table.
• from the bottom to top of the periodic table.
• High for the non metals (F is the highest)
• Low for the metals.

50
Q

Dipoles and bond polarity

A

A polar covalent bond becomes more polar as the difference in electronegativity increases.
The separation of charges (partial charges) in a polar bond is called a dipole.

51
Q

Polar molecules

A

Molecular polarity
• individual bond polarities and lone-pair electrons contribute
– e- are displaced toward the more electronegative atom.
• depends on the molecular shape of the compound
– Symmetrical molecules can have polar bonds and be nonpolar overall.

52
Q

Intermolecular Forces

A

Intermolecular forces determine the physical properties of molecular compounds.

• Mainly related to covalent molecules
– Ionic Compounds: substances made up of ions (not molecules) • Ionic bonds—strongest bond type (stronger than covalent bonds)
• Ionic and covalent bonds– intramolecular forces
• In gases, intermolecular forces are negligible.
• In liquids and solids, the stronger the intermolecular forces, the higher the melting and boiling points.
Three major types: dipole-dipole, London dispersion, hydrogen bonding.

53
Q

INTERmolecular forces

A

Forces between molecules
– Holds molecules (covalent compounds) together in solids and liquids
– Negligent in gases (molecules are too far apart for attractive forces to occur)
1. Dipole-Dipole attraction [polar molecules]
2. Hydrogen bonds (H-bonds) [polar molecules]
3. Dispersion/London Forces [main attraction btwn non-polar molecules and found between all other molecules]

54
Q

INTRAmolecular forces

A

Forces between atoms w/in a compound
– Holds atoms in a compound together
1. Ionic bonds (compounds not molecules; always solids @ Rm T due to strong electrostatic attraction)
2. Covalent bonds (molecules can be in either gases, liquids or solids @ Rm T)

55
Q

London dispersion forces
LDF

A

– All molecules experience these forces, but main attractive force in nonpolar
covalent molecules
– At any given instant there may be more electrons at one end of a molecule than at the other, giving a short-lived polarity.
– Weak attraction but make it possible for nonpolar molecules to form liquids and solids.
– The larger the molecular weight and surface area, the greater the temporary polarization of a molecule.

56
Q

Dipole-Dipole forces and Hydrogen Bonds

A

In polar covalent molecules..
• exert attractive forces called dipole-dipole attractions.
• form strong dipole attractions called hydrogen bonds
– Between molecules with H atoms bonded to either F, O, or N
• The H atom is attracted to the F, N, or O on another molecule
• In water, each oxygen atom has two lone pairs of e- and two hydrogen atoms, allowing the formation of four hydrogen bonds

57
Q

• Dipole-dipole forces:

A

– The positive and negative ends of different molecules are attracted to one another.
– The effect of dipole-dipole forces can be seen by the difference in boiling points between nonpolar and polar molecules.

58
Q

Comparison of bonding and attractive forces

A
59
Q

MP/BP and attractive forces

A

Melting points and boiling pts of compounds related to the strength of attractive forces between molecules or compounds.
• dispersion forces—>weak forces
• lower m.p. and b.p. (Covalent cmpds)
• hydrogen bonding—->stronger attractive forces • higher m.p. and b.p.
(Covalent cmpds)
• ionic compounds—>strong attractive forces between ions in the compound
• highest m.p. and b.p.

60
Q

Chemical formulas

A

-a shorthand way to represent chemical compounds.
-Chemical formulas represent each atom within the compound and the number of each atom by subscripts.
-The chemical formula of an ionic compound, is called a formula unit.
-It represents ratio of ions present in an ionic compound.

61
Q

molecular formula

A

-For covalent compounds the chemical formula is called a molecular formula
-represents the number and type of atoms present in one molecule.

62
Q

Ionic vs covalent bond table

A
63
Q

Naming compounds guide

A
64
Q

Common Polyatomic ion table

A
65
Q

1,2,3,4,5

A

Mono
Di
Tri
Tetra
Penta

66
Q

6,7,8,9,10

A

Hexa
Hepta
Octa
Nona
Deca

67
Q

Polyatomic ions

A

Polyatomic ions are groups of atoms covalently bonded to each other that have an unequal number of electrons and protons.
Most common polyatomic ions have a -1, -2, or -3 charge.
Two positively charged polyatomic ions are the hydronium ion, H3O+, and ammonia, NH4+.

68
Q

Ionic compounds or salts

A

-form in chemical reactions in which cations and anions chemically bind to each other by the transfer of electrons.
-Ionic compounds are formed between cations and anions bound by attractive electrostatic interactions.
-For an ionic solid, the interactions that hold the ions together make a crystalline lattice.
-Ions are firmly arranged in space through attractions between neighboring ions.
-This arrangement of ions minimizes space, and maximizes ionic bonding thereby increasing the stability of the compound.
-Most ionic compounds or salts, are brittle solids at room temperature with high melting and boiling points.
-some ionic compounds are soluble (can dissolve) in water
-the attraction between the water molecules that surround each ion in an ionic compound in solution is greater than the attractions between ions themselves in the ionic solid which is why an ionic compound can dissolve
-the ions are free to move about in the solution independent of each other and form an electrolyte solutions
-such a solution conducts electricity
-not all salts can dissolve in water

69
Q

Covalent compounds

A

-Unlike ions in salts (ionic compounds), the orbitals of each atom overlap to form covalent bonds.
-The energy required to completely separate electrons from each other is too great to overcome for non-metal atoms.
-usually, each atom contributes one electron to the bond.
-At other times one atom shares a pair of electrons with another atom in need of a pair of electrons; these are considered coordinate covalent bonds
-These types of bonds form between molecules such as ammonia and water with hydrogen ions (H+) to form the ammonium and hydronium cations, respectively.
-The more electrons shared between atoms, the stronger the bonds between them and the less freedom of rotation of atoms about the bonds.
-Nitrogen, oxygen, sulfur and phosphorus are common elements that are able to form double and triple bonds.
-Carbon dioxide and sulfur dioxide are binary molecules that both contain double bonds between atoms

70
Q

Seven diatomic molecules

A

H2, N2, O2, F2, Cl2, Br2, and I2
-have equal sharing of electrons bc atoms are pairing with identical atoms with equal electron affinities.
-The relative energetic stability of O2 and N2 molecules is only possible when oxygen atoms share 2 pairs (or 4) electrons while N atoms share 3 pairs (or 6) electrons instead of sharing only a single pair of electrons between atoms with the bonds in N2 being more stable than those in O2.

71
Q

Exceptions to octet rule

A

H following doubles rule as it has complete valence shell with 2 electrons
Boron has 3 valence electrons and usually forms 3 bonds with other atoms to have 6 valence electrons and forms stable compounds without octet of valence electrons
-Non-metals in the third period and below on the periodic table can form expanded octets by sharing unpaired d-orbital electrons and vacant d-orbitals allowing these atoms to make stable compounds in which they have more than an “octet” of valence electrons. Certain non-metals, like P, S, Cl, Br and I can form compounds with 10 or 12 valence electrons. Examples include SF6, BF3, and PF5.

72
Q

Polar covalent bonds & Electronegativity

A

-Not all covalent bonds form by an “equal” sharing of electrons between two atoms. The unequal sharing of electrons results in polar covalent bonds
-Atoms that are more likely to gain electrons from metal atoms have greater affinity for the electrons that they share with other atoms.
-This trend is called electronegativity and it mirrors the trends of ionization energy and electron affinity on the periodic table—
-increasing moving across from left to right in a period and decreasing moving down a group on the periodic table.
-Atoms of elements with high electronegativities will hold the electrons in a bond closer to their nuclei compared to atoms with relatively lower electronegativities.
-This unequal sharing can be represented by an arrow over the structure of the atoms, called a dipole, with a cross over the slightly positive atom and the point over the slightly negative atom

73
Q

[polyatomic ions]
Hydronium ion

A

H3O+

74
Q

[polyatomic ions]
Ammonium ion

A

NH4+

75
Q

[polyatomic ions]
Acetate Ion

A

CH3CO2 -

76
Q

[polyatomic ions]
Carbonate ion

A

CO3 2-

77
Q

[polyatomic ions]
Chlorite

A

ClO2 -

78
Q

[polyatomic ions]
Cyanide

A

CN-

79
Q

[polyatomic ions]
Dihydrogen phosphate ion

A

H2PO4 -

80
Q

[polyatomic ions]
Hydrogen carbonate ion
(Bicarbonate ion)

A

HCO3 -

81
Q

[polyatomic ions]
Hydrogen phosphate (biphosphate ion)

A

HPO4 2-

82
Q

[polyatomic ions]
Hydrogen sulfate ion (bisulfate)

A

HSO4 -

83
Q

[polyatomic ions]
Hydrogen sulfite (bisulfite)

A

HSO3 -

84
Q

[polyatomic ions]
Hydroxide

A

OH -

85
Q

[polyatomic ions]
Nitrate ion

A

NO3 -

86
Q

[polyatomic ions]
Nitrite

A

NO2 -

87
Q

[polyatomic ions]
Phosphate

A

PO4 3-

88
Q

[polyatomic ions]
Sulfate

A

SO4 2-

89
Q

[polyatomic ions]
Sulfite

A

SO3 2-

90
Q

[polyatomic ions]
Phosphite

A

PO3 3-

91
Q

[polyatomic ions]

A
92
Q

Methane
Carbon tetrahydride

A

CH4

Chem 100 Class (9 decks)

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