2
An octet is 8 valence electrons is associated with the stability of the noble gases does not occur with He; He is stable with 2 valence
electrons (duet)Valence Electrons
He 1s2 2
Ne 1s22s22p6 8
Ar 1s22s22p63s23p6 8
Kr 1s22s22p63s23p64s23d104p6 8
Octet Rule
3
Ionic and Covalent Bonds
Atoms form octets to become more stable by losing, gaining, or sharing
valence electrons by forming ionic or covalent
bonds
4
Metals Form Positive Ions
Metals form positive ions by a loss of their valence electrons with the electron configuration of the
nearest noble gas that have fewer electrons than protons
Group 1A(1) metals ion 1+
Group 2A(2) metals ion 2+
Group 3A(3) metals ion 3+
5
Formation of a Sodium Ion, Na+
Sodium achieves an octet by losing its one valence electron.
With the loss of its valenceelectron, a sodium ion has a 1+
charge. Sodium atom Sodium ion 11p+ 11p+
11e– 10e–
0 1 +
6
Formation of Magnesium Ion, Mg2+
Magnesium achieves an octet by losing its two valence electrons.
With the loss of two valenceelectrons magnesium forms apositive ion with a 2+ charge Mg atom Mg2+ ion 12p+ 12p+
12e– 10e–
0 2+
7
Formation of Negative Ions
In ionic compounds, nonmetals achieve an octet arrangement gain electrons form negatively charged ions with
3–, 2–, or 1– charges
8
Formation of a Chloride Ion, Cl–
Chlorine achieves an octet by adding an electron toits valence electrons.
9
Ionic Charge from Group Numbers
The charge of a positive ion is equal to its Group number.Group 1A(1) = 1+
Group 2A(2) = 2+
Group 3A(3) = 3+
The charge of a negative ion is obtained by subtracting 8 or 18 from its Group number .
Group 6A(16) = 6 – 8 = 2–
or 16 – 18 = 2–
Group Number and Ionic Charge
Ions achieve the electron configuration of their nearest noble gas of metals in Groups 1A(1), 2A(2), or 3A(13) have positive
1+, 2+, or 3+ charge. Of nonmetals in Groups 5A(15), 6A(16), or 7A(17) have
negative 3–, 2–, or 1– charge.
10
13
Ionic compounds consist of positive and negative ions have attractions called ionic bonds between positively
and negatively charged ions have high melting and boiling points are solid at room temperature
Ionic Compounds
15
An ionic formula consists of positively and negatively charged ions is neutral has charge balance
total positive charge = total negative charge
The symbol of the metal is written first, followed by the symbol of the nonmetal.
Ionic Formulas
16
Charge Balance for NaCl, “Salt”In NaCl, a Na atom loses its valence electron a Cl atom gains an electron the symbol of the metal is written first, followed by the
symbol of the nonmetal.
17
Charge Balance In MgCl2In MgCl2, a Mg atom loses two valence electrons two Cl atoms each gain one electron subscripts indicate the number of ions needed to give
charge balance
18
Writing Ionic Formulas from Charges
Charge balance is used to write the formula forsodium nitride, a compound containing Na+ and N3−.
Na+
3 Na+ + N3− = Na3N
Na+
3(1+) + 1(3–) = 0
20
Naming Ionic Compounds with Two Elements
To name a compound with two elements, identify the cation and
anion name the cation first,
followed by the name of the anion
21
Formula Ions Name Cation Anion
NaCl Na+ Cl– sodium chloride
K2S K+ S2– potassium sulfideMgO Mg2+ O2– magnesium oxide
CaI2 Ca2+ I– calcium iodide
Al2O3 Al3+ S2– aluminum sulfide
Examples of Ionic Compounds with Two Elements
22
Transition Metals Form Positive Ions
Most transition metals and Group 4(14) metals, Form 2 or more positive ions Zn2+, Ag+, and Cd2+ form only one ion.
Metals with Variable Charge
The names of transitionmetals with two or morepositive ions (cations) use aRoman numeral after thename of the metal to identifyionic charge.
23
25
Naming FeCl2
STEP 1 Determine the charge of the cation from the anion.
Fe ion + 2 Cl– = Fe ion + 2– = 0 Fe ion = 2+ = Fe2+
STEP 2 Name the cation by the element name, and use a Roman numeral to show its charge. Fe2+ = iron(II)
STEP 3 Write the anion with an ide ending. chlorideSTEP 4 Name the cation first, then the anion. iron(II) chloride
26
Naming Cr2O3
STEP 1 Determine the charge of cation from the anion. 2Cr ions + 3O2– = 2Cr ions + 3(2–)
= 2Cr ions + 6– = 0 2Cr ions = 6+ Cr ion = 3+ = Cr3+
STEP 2 Name the cation by the element name, and use a Roman numeral to show its charge. Cr3+ = chromium(III)
STEP 3 Write the anion with an ide ending. oxideSTEP 4 Name the cation first, then the anion. chromium (III) oxide
28
Writing Formulas
Write a formula for potassium sulfide.STEP 1 Identify the cation and anion. potassium = K+
sulfide = S2−
STEP 2 Balance the charges. K+ S2−
K+
2(1+) + 1(2–) = 0STEP 3 Write the cation first.
2K+ and 1S2− = K2S1 = K2S
29
Writing Formulas
Write a formula for iron(III) chloride.
STEP 1 Identify the cation and anion. iron (III) = Fe3+ (III = charge of 3+) chloride = Cl−
STEP 2 Balance the charges. Fe3+ Cl−
Cl− Cl−
1(3+) + 3(1–) = 0
STEP 3 Write the cation first.1Fe3+ and 3Cl− = FeCl3
31
A polyatomic ion is a group of atoms has an overall ionic chargeExamples:
NH4+ ammonium OH− hydroxide
NO3−
nitrate NO2−
nitrite
CO32− carbonate PO4
3− phosphate
HCO3− hydrogen carbonate
(bicarbonate)
Polyatomic Ions
32
The names of common polyatomic anions end in ate
NO3− nitrate PO4
3− phosphate with one oxygen less end in ite
NO2− nitrite PO3
3− phosphite with hydrogen attached use prefix hydrogen (or bi)
HCO3− hydrogen carbonate (bicarbonate)
HSO3− hydrogen sulfite (bisulfite)
Some Names of Polyatomic Ions
34
The positive ion is named first followed by the name of the polyatomic ion.
NaNO3 sodium nitrate
K2SO4 potassium sulfate
Fe(HCO3)3 iron(III) bicarbonate
or iron(III) hydrogen carbonate (NH4)3PO3 ammonium phosphite
Examples of Names of Compounds with Polyatomic Ions
35
Writing Formulas with Polyatomic Ions
The formula of an ionic compound containing a polyatomic ion must have a charge balance
that equals zero (0)Na+ and NO3
− NaNO3
with two or more polyatomic ions put the polyatomic ions in parentheses.
Mg2+ and 2NO3−
Mg(NO3)2
subscript 2 for charge balance
37
Forming Octets in Molecules
In a fluorine (F2) molecule, each F atom shares one electron acquires an octet
40
Single and Multiple Bonds
In a single bond, one pair of electrons is shared.
In a double bond, two pairs of electrons are shared.
In a triple bond, three pairs of electrons are shared.
41
Electron-Dot Formula of CS2
Write the electron-dot formula for CS2.STEP 1 Determine the atom arrangement. The C atom is the
central atom.
S C SSTEP 2 Determine the total number of valence electrons for 1C
and 2S. 1 C(4e–) + 2 S(6e–) = 16e–
STEP 3 Attach each S atom to the central C atom using one electron pair.
S : C : S 16e– – 4e– = 12e– remaining
STEP 4 Attach 12 electrons as 6 lone pairs. .. .. : S : C : S :
42
Electron-Dot Formula of CS2 (continued)
To complete octets, form one or more multiple bonds. Convert two lone pairs to bonding pairs between C and S atoms to make two double bonds.
43
A Nitrogen Molecule has A Triple Bond
In a nitrogen molecule, N2, each N atom shares 3 electrons each N atom attains an octet the sharing of 3 sets of electrons is a multiple bond
called a triple bond
44
Resonance structures are two or more electron-dot formulas for the same
arrangement of atoms related by a double-headed arrow ( ) written by changing the location of a double bond
between the central atom and a different attached atom
Resonance Structures
45
Sulfur dioxide has two resonance structures. STEP 1 Write the arrangement of atoms.
O S O STEP 2 Determine the total number of valence electrons. 1 S(6e−) + 2 O(6e−) = 18e−
STEP 3 Connect bonded atoms by single electron pairs.
O : S : O 4e− used 18e− – 4e− = 14e− remaining
Writing Resonance Structures
46
STEP 4 Add 14 remaining electrons as 7 lone pairs.
STEP 5 Form a double bond to complete octets. Two resonance structures are possible.
Writing Resonance Structures (continued)
47
Write the ionic formula of the compound containing Ba2+ and Cl.
Write the symbols of the ions.Ba2+ Cl
Balance the charges. Ba2+ Cl two Cl needed
Cl
Write the ionic formula using a subscript 2 for two chloride ions that give charge balance.
BaCl2
Formula from Ionic Charges
48
Chapter 5Compounds and Their Bonds
5.6Naming and Writing Covalent Formulas
NO nitrogen oxideNO2 nitrogen dioxide
N2O4 dinitrogen tetroxide
49
Names of Covalent Compounds
Prefixes are used in the names of covalent compounds because two nonmetals can form two or more different
compoundsExamples of compounds of N and O:
NO nitrogen oxideNO2 nitrogen dioxide
N2O dinitrogen oxide
N2O4 dinitrogen tetroxide
N2O5 dinitrogen pentoxide
50
Naming Covalent Compounds
STEP 1 Name the first nonmetal by its element name. STEP 2 Name the second nonmetal with an ide ending.STEP 3 prefixes to indicate the number (from
subscripts) of atoms of each nonmetal. Mono is usually omitted.
51
What is the name of SO3?
STEP 1 The first nonmetal is S sulfur. STEP 2 The second nonmetal is O, named oxide.STEP 3 The subscript 3 of O is shown as the prefix tri.
SO3 → sulfur trioxide
The subscript 1(for S) or mono is understood.
Naming Covalent Compounds (Continued)
52
Name P4S3
STEP 1 The first nonmetal P is phosphorus. STEP 2 The second nonmetal S is sulfide.STEP 3 The subscript 4 of P is shown as tetra.
The subscript 3 of O is shown as tri.
P4S3 → tetraphosphorus trisulfide
Naming Covalent Compounds (Continued)
54
Write the formula for carbon disulfide.
STEP 1 Elements are C and SSTEP 2 No prefix for carbon means 1 C
Prefix di = 2 Formula: CS2
Writing Formulas of Covalent Compounds
56
The electronegativity value indicates the attraction of an atom for shared electrons increases from left to right going across a period on the
periodic table decreases going down a group on the periodic table is high for the nonmetals, with fluorine as the highest is low for the metals
Electronegativity
58
A nonpolar covalent bond occurs between nonmetals has an equal or almost equal sharing of electrons has almost no electronegativity difference (0.0 to 0.4)
Examples: Atoms Electronegativity Type of Bond
Difference _______________N–N 3.0 – 3.0 = 0.0 Nonpolar covalentCl–Br 3.0 – 2.8 = 0.2 Nonpolar covalentH–Si 2.1 – 1.8 = 0.3 Nonpolar covalent
Nonpolar Covalent Bonds
59
A polar covalent bond occurs between nonmetal atoms has an unequal sharing of electrons has a moderate electronegativity difference
(0.5 to 1.7)
Examples: Atoms Electronegativity Type of Bond Difference_________________________________________
O–Cl 3.5 – 3.0 = 0.5 Polar covalentCl–C 3.0 – 2.5 = 0.5 Polar covalentO–S 3.5 – 2.5 = 1.0 Polar covalent
Polar Covalent Bonds
62
VSEPR
In the valence-shell electron-pair repulsion theory(VSEPR), the electron groups around a central atom are arranged as far apart from each other as possible have the least amount of repulsion of the negatively
charged electrons have a geometry around the central atom that
determines molecular shape
63
Shapes of Molecules
The three-dimensional shape of a molecule is the result of bonded groups and lone pairs of electrons
around the central atom is predicted using the VSEPR theory (valence-shell-
electron-pair repulsion)
65
Two Electron Groups
In a molecule of BeCl2, there are two electron groups bonded to the central
atom, Be (Be is an exception to the octet rule) .. .. : Cl : Be : Cl :
.. ..
to minimize repulsion, the arrangement of two electron groups is 180°, or opposite each other
the shape of the molecule is linear
66
Two Electron Groups with Double Bonds
In a molecule of CO2, there are two electron groups
bonded to C (electrons in each double bond are counted as one group)
repulsion is minimized with the double bonds opposite each other at 180°
the shape of the molecule is linear
67
Three Electron Groups
In a molecule of BF3, three electron groups are bonded to
the central atom B (B is an exception to the octet rule)
.. : F:
.. .. .. : F : B : F :
.. .. repulsion is minimized with 3 electron
groups at angles of 120° the shape is trigonal planar
68
Two Electron Groups and One Lone Pair
In a molecule of SO2, S has 3 electron groups; 2 electron
groups bonded to O atoms and one lone pair .. .. ..
:O :: S : O : ..
repulsion is minimized with the electron groups at angles of 120°, a trigonal planar arrangement
the shape is bent (120°), with two O atoms bonded to S
● ●
69
Four Electron Groups
In a molecule of CH4, there are four electron groups around C repulsion is minimized by placing four electron groups at
angles of 109°, which is a tetrahedral arrangement the four bonded atoms form a tetrahedral shape
70
Three Bonding Atoms and One Lone Pair
In a molecule of NH3, three electron groups bond to H atoms, and the fourth one
is a lone (nonbonding) pair repulsion is minimized with 4 electron groups in a
tetrahedral arrangement the three bonded atoms form a pyramidal (~109°) shape
71
Two Bonding Atoms and Two Lone Pairs
In a molecule of H2O, two electron groups are bonded to H atoms and two are
lone pairs (4 electron groups) four electron groups minimize repulsion in a tetrahedral
arrangement the shape with two bonded atoms is bent (~109°)
Determining Molecular Polarity
Determine the polarity of the H2O molecule.
Solution: The four electron groups of oxygen are bonded to two H atoms. Thus the H2O molecule has a net dipole, which makes it a polar molecule.
72
74
Dipole–Dipole Attractions
In covalent compounds, polar molecules exert attractive forces called dipole-dipole
attractions form strong dipole attractions called hydrogen
bonds between hydrogen atoms bonded to F, O, or N, and other atoms that are very electronegative
76
Dispersion Forces
Dispersion forces are weak attractions between nonpolar molecules caused by temporary dipoles that develop when
electrons are not distributed equally
78
Melting Points and Attractive Forces
Ionic compounds require large amounts of energy to break apart ionic bonds. Thus, they have high melting points.
Hydrogen bonds are the strongest type of dipole–dipole attractions. They require more energy to break than do other dipole attractions.
Dispersion forces are weak interactions, and very little energy is needed to change state.