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Bonding
• Please read and answer the sheet you have been given.
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Energy and Chemical Bonds• Chemical bonds are the forces that hold atoms
together in a compound.
• Energy is required to overcome these attractive forces and separate the atoms in a compound. Thus, the breaking of a chemical bond is an endothermic process.
• If energy is required to break a bond, then the opposite process of forming a bond must release energy. The formation of a bond is an exothermic process.
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Question
• When a chemical bond is broken, the resulting compound has more potential energy than the substance from which it was formed. Why?– Breaking bonds requires energy, therefore the new
compound will have more energy than at the beginning.
Endothermic
AB + energy → A + B
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Question
• Conversely, when a chemical bond is formed, the resulting compound has less potential energy than the substances from which it was formed. Why?– Forming bonds releases energy, therefore the new
compound will have less energy than at the beginning.
Exothermic
A + B → AB + energy4
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Chemical BondingRECAP
Page 4Complete top half of page.
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How can we make Na and Cl happy?
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Endothermic or Exothermic?
Exothermic
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Chemical BondingRECAP
Page 4Complete bottom half of page.
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1.) What is a 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
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MolecularFormula
FormulaUnit
IONIC COVALENT
COCO22NaClNaCl
CHEMICAL FORMULA
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COMPOUND
TernaryCompound
BinaryCompound
2 elementsmore than 2
elements
NaNONaNO33NaClNaCl
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ION
PolyatomicIon
MonatomicIon
1 atom 2 or more atoms
NONO33--NaNa++
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Chemical bonds are formed when valence electrons are:
• transferred from one atom to another (ionic)
• shared between atoms (covalent)• mobile within a metal (metallic)
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Ionic bonds are formed when metals transfer their valence electrons to nonmetals.The oppositely charged ions attract each other to form an ionic bond.
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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
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Electron-dot diagrams (Lewis structures) can represent the valence electron arrangement in elements, compounds, and ions.
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atom ion molecular compound
ionic compound
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Dots represent valence electrons.Everything else (inner shell electrons and nucleus) is called the Kernel and
is represented by the symbol.
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Phosphorous has 5 valence electrons so we draw 5 dots around the symbol for phosphorous.
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Draw the Lewis Dot Structures of the first 18 elements.
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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.
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Magnesium atom Magnesium ion
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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.
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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.
23The formula for magnesium fluoride is MgF2
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Two major categories of compounds are ionic and
molecular (covalent) compounds.
• 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.
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Comparing the properties compounds with ionic bonds and compounds with covalent bonds.
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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
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Ionic solids conduct electricity when dissolved or molten.
Molecular solids do not.
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Ionic Solid dissolved in water
Molecular Solid dissolved in water
Solution conducts electricity
Solution doesn’t conduct
electricity
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Nomenclature
“Or How Do We Name Compounds”
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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
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Anion – Negative ion– Has gained electrons– Non metals form
anions
Cation– Positive ion– Formed by losing
electrons– Metals form cations
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Ionic Compounds
• Made of cations and anions• Metals and nonmetals• Electrons lost by the cation are gained by the
anion
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Ionic Compounds
Na + Cl
Sodium is cation
1-
ClNa +1+
Chlorine is anion
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Charges on Ions
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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
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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
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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+
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Naming Anions
• Anions are always the same.• Change the element ending to -- ide• F1- Fluorine to Fluoride
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Name These
• Cl1-
• N3-
• Br 1-
• O2-
• I1-
• Sr2+
• Chloride• Nitride• Bromide• Oxide• Iodide• Strontium
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Write These
• Sulfide ion• Iodide ion• Phosphide ion• Strontium ion
• S2-
• I1-
• P3-
• Sr2+
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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-
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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
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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.
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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
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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
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Polyatomic ions are groups of atoms covalently bonded
together that have a negative or positive charge.
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Polyatomic ions are held together by covalent bonds but
form ionic bonds with other ions.
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H N H Cl
H
H
+
-Covalent bonds
Ionic bond
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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
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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
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Writing Formulas• Crisscross method
Ca2+ Cl1- CaCl2No need to write the oneIron (III) sulfide
Calcium chloride
Fe 2 S3
Fe 3+ S2-
Fe2S3
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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
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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
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COVALENT BONDbond formed by the sharing of electrons
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Covalent Bond
• Between nonmetallic elements of similar electronegativity.
• Formed by sharing electron pairs• Stable non-ionizing particles, they are not
conductors at any state• Examples; O2, CO2, C2H6, H2O, SiC
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Bonds in all the polyatomic ions and diatomics
are all covalent bonds
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when electrons are shared equally
NONPOLAR COVALENT BONDS
H2 or Cl2
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2. Covalent bonds- Two atoms share one or more pairs of outer-shell electrons.
Oxygen AtomOxygen Atom Oxygen AtomOxygen Atom
Oxygen Molecule (O2)
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when electrons are shared but shared
unequally
POLAR COVALENT BONDS
H2O
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Polar Covalent Bonds: Unevenly matched, but willing to share.
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- water is a polar molecule because oxygen is more electronegative than hydrogen, and therefore electrons are pulled closer to oxygen.
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when electrons are from one element
only
COORDINATE COVALENT BONDS
NH4+1 or H3O+1
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Homework Page 5
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Molecular Compounds
Writing Names and Formulas
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Covalent Bonding / Compounds
• Compounds in which the electronegativity difference is less than 1.7
• 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
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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
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In a multiple covalent bond, more than one pair of electrons are shared between two atoms.
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•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
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Regents Question: 08/02 #17
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Which molecule contains a triple covalent
bond?
(1) H 2
(2) N 2
(3) O 2
(4) Cl 2
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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)
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A polar molecule is called a dipole. It has a positive side and a negative side – uneven charge distribution.
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Symmetrical (nonpolar) molecules include CO2 ,
CH4 , and diatomic elements. ..
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Symmetrical molecules are not dipoles.
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Asymmetrical (polar) molecules include HCl, NH3 , and H2 O. (5.2l)
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The negative side of the molecule is the side that has the atom with the higher electronegativity.
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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
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• Homework Page 10
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Two pages after 11 before 12
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Page 13 for notes
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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.
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Naming
• The second part of all names end with -ide
• Prefixes are used to indicate number of each atom
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Prefixes
• 1 mono-• 2 di-• 3 tri-• 4 tetra-• 5 penta-• 6 hexa-• 7 hepta-• 8 octa-
• 9 nona-• 10 deca-
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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)
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Name These
• N2O
• NO2
• Cl2O7
• CBr4
• CO2
• BaCl2
• H2O
• Dinitrogen monoxide• Nitrogen dioxide• Dichlorine heptoxide• Carbon tetrabromide• Carbon dioxide• Barium chloride• Dihydrogen monoxide
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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
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Page 13 practice
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Page 21 summary
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ionic covalent
valence electrons
Comparison of Bonding Types
sharing of electrons
transfer of electrons
ionsmolecules
EN > 1.7 EN < 1.7
high mp low mp
molten salts conductive
non-conductive
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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 88
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METALLIC BONDbond found in
metals; holds metal atoms together
very strongly
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Metallic Bond
• Formed between atoms of metallic elements• Electron cloud around atoms • Good conductors at all states, lustrous, very
high melting points• Examples; Na, Fe, Al, Au, Co
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Metallic Bonds: Mellow dogs with plenty of bones to go around.
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Metallic Bond, A Sea of Electrons
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Metals Form Alloys
Metals do not combine with metals. They form Alloys which is a solution of a metal in a metal.Examples are steel, brass, bronze and pewter.
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Intermolecular Forces
• Weaker than covalent bonds• Weak intermolecular forces – lower boiling point
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The stronger the intermolecular forces, the higher the boiling
points and melting points.
• Ionic Solids• Molecules with Hydrogen bonds• Polar molecules• Nonpolar molecules
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Strongest
Weakest
For nonpolar molecules, the greater the mass, the greater the force of attraction.
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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
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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.
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Hydrogen Bonds
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Hydrogen Bonds
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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
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ion-dipole forces
• Attractive forces between neutral molecules and charged (ionic) compounds
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Ion-dipole forces(Ion-Molecule attraction)
•are important in solutions of ionic substances in polar solvents •(e.g. a salt in aqueous solvent)