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Page 2. Chemical Bonding. Chemical Bond attractive force between atoms or ions that binds them together as a unit bonds form in order to… decrease potential energy (PE) increase stability. CHEMICAL FORMULA. IONIC. COVALENT. Formula Unit. Molecular Formula. NaCl. CO 2. COMPOUND. - PowerPoint PPT Presentation
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Chemical Bonding
Page 2
• Chemical Bond– attractive force between atoms or ions that binds
them together as a unit
– bonds form in order to…• decrease potential energy (PE)• increase stability
MolecularFormula
FormulaUnit
IONIC COVALENT
COCO22NaClNaCl
CHEMICAL FORMULA
COMPOUND
TernaryCompound
BinaryCompound
2 elements more than 2elements
NaNONaNO33NaClNaCl
ION
PolyatomicIon
MonatomicIon
1 atom 2 or more atoms
NONO33--NaNa++
Chemical bonds are formed when valence electrons are:• transferred from one atom to another
(ionic)• shared between atoms (covalent)• mobile within a metal (metallic)
6
Ionic bonds are formed when metals transfer their valence electrons to nonmetals.The oppositely charged ions attract each other to form an ionic bond.
7
Sodium has one valence electron and chlorine has seven. Sodium want to lose 1 electron and chlorine needs to gain 1.
Sodium transfers its valence electron to chlorine
Forming an Na+ and a Cl- ion – sodium chloride NaCl
Electron-dot diagrams (Lewis structures) can represent the valence electron arrangement in elements, compounds, and ions.
8
atom ion molecular compound
ionic compound
Dots represent valence electrons.Everything else (inner shell electrons and nucleus) is called the Kernel and
is represented by the symbol.
9
Phosphorous has 5 valence electrons so we draw 5 dots around the symbol for phosphorous.
Draw the Lewis Dot Structures of the first 18 elements.
10
When metals lose electrons to form ions, they lose all their
valence electrons. The Lewis Dot Structure of a metal ion has no dots. The charge indicates how
many electrons were lost.
11
Magnesium atom Magnesium ion
When nonmetals gain electrons, they fill up their valence shell with a complete octet (except hydrogen.) The ion is placed in
brackets with the charge outside the brackets.
12
A + metal ion is attracted to a – nonmetal ion (opposites attract)
forming an ionic compound. We can use Lewis dot structures to represent
ionic compounds.
13The formula for magnesium fluoride is MgF2
Two major categories of compounds are ionic and
molecular (covalent) compounds. (5.2g)
• Ionic compounds are formed when a metal combines with a nonmetal.
• Ionic compounds have ionic bonds.
• Molecular compounds are formed between two nonmetals.
• Molecular compounds have covalent bonds.
14
Comparing the properties compounds with ionic bonds and compounds with covalent bonds.
15
Properties of ionic compounds– Solids with high melting
and boiling points (strong attraction between ions)
– Electrolytes: Do not conduct electricity as solids but do when dissolved or molten – ions are charged particles that are free to move
– No individual molecules
Properties of molecular compounds– Low melting and boiling
points (weak attraction between molecules)
– Nonelectrolytes: Do not conduct electricity as solids or when dissolved or molten – no charged particles (ions) to move
– Solids are soft– Forms molecules
Ionic solids conduct electricity when dissolved or molten.
Molecular solids do not.
16
Ionic Solid dissolved in water
Molecular Solid dissolved in water
Solution conducts electricity
Solution doesn’t conduct
electricity
Nomenclature
“Or How Do We Name Compounds”
Systematic Naming
• Compound is made up of two or more elements
• Name should tell us how many and what type of atoms
• Too many compounds to remember all the names
Anion – Negative ion– Has gained electrons– Non metals form
anions
Cation– Positive ion– Formed by losing
electrons– Metals form cations
Ionic Compounds
• Made of cations and anions• Metals and nonmetals• Electrons lost by the cation are gained by the
anion
Ionic Compounds
Na + Cl
Sodium is cation
1-
ClNa +1+
Chlorine is anion
Charges on Ions
Naming Ions
• Metal ion is written first in both name and formula– It is named directly from element which formed the ion.– Will nearly always be the positive ion or “cation”
– Transition metals can have more than one type of charge– Indicate the charge with roman numerals in parenthesis.
Iron(II) or Iron(III) – Exceptions:
• Silver always +1 • Cadmium and Zinc always +2
Name these
• Na 1+
• Ca 2+
• Al 3+
• Fe 3+
• Fe 2+
• Pb 2+
• Li 1+
• Sodium• Calcium• Aluminum• Iron (III)• Iron (II)• Lead (II)• Lithium
Write Formulas for these
• Potassium ion• Magnesium ion• Copper (II) ion• Chromium (VI) ion • Barium ion• Mercury (II) ion
• K1+
• Mg2+
• Cu2+
• Cr6+
• Ba2+
• Hg2+
Naming Anions
• Anions are always the same.• Change the element ending to -- ide• F1- Fluorine to Fluoride
Name These
• Cl1-
• N3-
• Br 1-
• O2-
• I1-
• Sr2+
• Chloride• Nitride• Bromide• Oxide• Iodide• Strontium
Write These
• Sulfide ion• Iodide ion• Phosphide ion• Strontium ion
• S2-
• I1-
• P3-
• Sr2+
Polyatomic Ions• Tightly bound groups of atoms acting as a
single ion.• Names given in table in book. (pg 123)• Most are anions that contain oxygen. Names
end in –ate (one more O), or –ite (one less O).• SO3
2- = sulfite; SO42- = sulfate
• Exceptions: Ammonium cation NH4+, Cyanide CN-, and hydroxide OH-
Naming Binary Ionic Compounds
• 2 elements involved• Ionic – metal (cation) and a non-metal (anion)• Naming is easy with representative elements
in A groups• NaCl = Na+ Cl- = sodium chloride• MgBr2 = Mg2+Br- = magnesium bromide
Naming Binary Ionic Compounds
• The problem comes with the transition metals.
• Need to figure out their charges• All ionic compounds will have a neutral charge– Same number of + and – charges
• Use the anion to determine the charge on the positive ion.
Naming Binary Ionic Compounds
• Try naming these– KCl– Na3N– CrN– ScP– PbO– PbO2
– Na2Se
– Potassium chloride– Sodium nitride– Chromium (III) nitride– Scandium (III) phosphide– Lead (II) oxide– Lead (IV) oxide– Sodium selenide
Tertiary Ionic Compounds• Will have polyatomic ions• At least 3 elements• Use blue sheet• Name these ions– NaNO3
– CaSO4
– CuSO3
– (NH4)2O
– LiCN– Fe(OH)3
– (NH4)2CO3
– NiPO4
•Sodium nitrate
•Calcium sulfate
•Copper (II) sulfite
•Ammonium oxide
• Lithium cyanide• Iron (III) hydroxide• Ammonium carbonate• Nickel (III) phosphate
Polyatomic ions are groups of atoms covalently bonded
together that have a negative or positive charge.
34
Polyatomic ions are held together by covalent bonds but
form ionic bonds with other ions.
35
H N H Cl
H
H
+-Covalent
bonds
Ionic bond
Writing Formulas
• Charges have to add up to zero.• Get charges on pieces from Periodic Table• Cations from element name on table• Anions from table change ending to –ide, or
use name of polyatomic ion• Balance the charges • Put polyatomics in parenthesis
Writing Formulas
• Write formula for calcium chloride– Calcium is Ca2+
– Chloride is Cl1-
– Ca+2Cl-1 would have a +1 charge– Need another Cl1-
– Ca+2Cl2-1 = CaCl2
Writing Formulas• Crisscross method
Ca2+ Cl1- CaCl2No need to write the oneIron (III) sulfide
Calcium chloride
Fe 2 S3
Fe 3+ S2-
Fe2S3
Write Formulas for These• Lithium sulfide• Tin (II) oxide• Tin (IV) oxide• Magnesium fluoride• Copper (II) sulfate• Iron (III) phosphide• Iron (III) sulfide• Ammonium chloride• Ammonium sulfide
• Li2S• SnO• SnO2
• MgF2
• CuSO4
• FeP• Fe2S3
• (NH4)Cl• (NH4)2S
Things to Look For
• If cations have ( ), the roman numeral is their charge.
• If anions end in –ide they probably are off the periodic table (monoatomic)
• If anion ends in –ate or –ite it is a polyatomic ion
Molecular Compounds
Writing Names and Formulas
Covalent Bonding / Compounds
• Compounds in which the electronegativity difference is less than 2.0
• Between a nonmetal and nonmetal• Can’t be held together because of opposite
charges• Can’t use charges to figure out how many of
each atom
Covalent Bonding
• Smallest piece of a covalently bonded compound is a molecule
• Electrons are shared between atoms in bond
Water
H2O
Carbon Dioxide
CO2 Ammonia
NH3
In a multiple covalent bond, more than one pair of electrons are shared
between two atoms. (5.2e)
44
•Diatomic oxygen has a double bond O=O (2 shared pairs) because oxygen needs 2 electrons to fill its valence shell
•Diatomic nitrogen has a triple bond NN (3 shared pairs) because nitrogen needs 3 electrons to fill its valence shell
•Carbon dioxide has two double bonds
Regents Question: 08/02 #17
45
Which molecule contains a triple covalentbond?(1) H 2
(2) N 2
(3) O 2
(4) Cl 2
Molecular polarity can be determined by the shape of the molecule and the
distribution of charge.• Possible shapes– Linear (X2 HX CO2)
– Bent (H2O)
– Pyramidal (NH3)
– Tetrahedral (CH4 CCl4)
46
A polar molecule is called a dipole. It has a positive side and a negative side – uneven charge distribution.
Symmetrical (nonpolar) molecules include CO2 ,
CH4 , and diatomic elements. ..
47
Symmetrical molecules are not dipoles.
Asymmetrical (polar) molecules include HCl, NH3 , and H2 O. (5.2l)
48
The negative side of the molecule is the side that has the atom with the higher electronegativity.
Differences between ionic and covalent bonding:
Na + Cl-
ClNa ++
Ionic bonding• electron is “stolen”
• high electronegativity difference
• between metal & nonmetal
• Formation of crystal structure
think proportions of atoms in
formula unit NaCl 1:1
Molecules are easier to name and work with
• Ionic compounds use charges to determine how many of each.– Have to figure out charges– Have to figure out numbers
• Molecular compound’s name tells you the number of atoms.
Naming
• The second part of all names end with -ide
• Prefixes are used to indicate number of each atom
Prefixes
• 1 mono-• 2 di-• 3 tri-• 4 tetra-• 5 penta-• 6 hexa-• 7 hepta-• 8 octa-
• 9 nona-• 10 deca-
Naming Continued
• To write the name…write two wordsPrefix-name Prefix-name –ide
• One exception is we don’t write mono- if there is only one of the first element.
• No double vowels when writing names– (oa oo)
Name These
• N2O
• NO2
• Cl2O7
• CBr4
• CO2
• BaCl2
• H2O
• Dinitrogen monoxide• Nitrogen dioxide• Dichlorine heptoxide• Carbon tetrabromide• Carbon dioxide• Barium chloride• Dihydrogen monoxide
Write Formulas for These• Diphosphorous pentoxide• Tetraiodine monoxide• Sulfur hexaflouride• Nitrogen trioxide• Carbon tetrahydride• Phosphorous trifluoride• Aluminum chloride
• P2O5
• I4O
• SF6
• NO3
• CH4
• PFl3
• AlCl3
Lewis Dot Structure(AKA Electron Dot Structure)
1. Write the symbol for each atom and show each of their valence electrons as dots (ignore all electrons below valence shell)
Cl Cl
Cl
Cl Cl
2. The number of electrons before you combine the atoms will equal number you have after.
Cl2
ionic covalent
valence electrons
Comparison of Bonding Types
sharing of electrons
transfer of electrons
ions molecules
EN > 1.7 EN < 1.7
high mp low mp
molten salts conductive
non-conductive
The bonds holding metals together in their crystal lattice
are called metallic bonds.
• All metals have metallic bonds• “Positive ions immersed in a sea of mobile
electrons”– Bonds are between Kernels, leaving the valence
electrons free to move from atom to atom– Mobile electrons give metals the ability to
conduct electricity 58
Intermolecular Forces• Weaker than covalent bonds• Weak intermolecular forces – lower boiling point
The stronger the intermolecular forces, the higher the boiling
points and melting points.
• Ionic Solids• Molecules with Hydrogen bonds• Polar molecules• Nonpolar molecules
61
Strongest
Weakest
For nonpolar molecules, the greater the mass, the greater the force of attraction.
Hydrogen Bonds
• Hydrogen bonds are considered to be dipole-dipole type interactions
• Hydrogen bonds vary from about 4 kJ/mol to 25 kJ/mol (so they are still weaker than typical covalent bonds.
• But they are stronger than dipole-dipole and or dispersion forces.
Hydrogen Bonds
Hydrogen Bonds
ion-dipole forces
• Attractive forces between neutral molecules and charged (ionic) compounds
Ion-dipole forces(Ion-Molecule attraction)
•are important in solutions of ionic substances in polar solvents •(e.g. a salt in aqueous solvent)
Van der Waals Forces
• Weak bonds• Liquefy gases• Bonds that combine gas molecules to form liquid• Ex. CO2 – liquid in toy car
- liquid nitrogen• Molecules must be close to each other• Larger atoms have stronger Van-der Waals forces
Dipole-dipole Forces• Polar molecules attract one another when the partial positive
charge on one molecule is near the partial negative charge on the other molecule
• The polar molecules must be in close proximity for the dipole-dipole forces to be significant
• Dipole-dipole forces are characteristically weaker than ion-dipole forces
• Dipole-dipole forces increase with an increase in the polarity of the molecule
London Dispersion Forces • Nonpolar molecules would not seem to have any basis for
attractive interactions • However, gases of nonpolar molecules can be liquefied
indicating that if the kinetic energy is reduced, some type of attractive force can predominate.
• Fritz London (1930) suggested that the motion of electrons within an atom or non-polar molecule can result in a transient dipole moment
London Forces