Chemical Bonding Page 2 Chemical Bond attractive force

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

CHEMICAL FORMULA IONIC COVALENT Formula Unit Molecular Formula Na. Cl CO 2

COMPOUND 2 elements Binary Compound Na. Cl more than 2 elements Ternary Compound Na. NO 3

ION 1 atom Monatomic Ion + Na 2 or more atoms Polyatomic Ion NO 3

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. 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 Na. Cl 7

Electron-dot diagrams (Lewis structures) can represent the valence electron arrangement in elements, compounds, and ions. atom ion molecular compound ionic compound 8

Dots represent valence electrons. Everything else (inner shell electrons and nucleus) is called the Kernel and is represented by the symbol. Phosphorous has 5 valence electrons so we draw 5 dots around the symbol for phosphorous. 9

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. Magnesium atom Magnesium ion 11

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. The formula for magnesium fluoride is Mg. F 2 13

Two major categories of compounds are ionic and molecular (covalent) compounds. (5. 2 g) • 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. ¨ 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 15 – Forms molecules

Ionic solids conduct electricity when dissolved or molten. Molecular solids do not. Solution doesn’t conduct electricity Solution conducts electricity Ionic Solid dissolved in water Molecular Solid dissolved in water 16

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

Cation – Positive ion – Formed by losing electrons – Metals form cations Anion – Negative ion – Has gained electrons – Non metals form anions

Ionic Compounds • Made of cations and anions • Metals and nonmetals • Electrons lost by the cation are gained by the anion

Ionic Compounds Sodium is cation Na + Cl 1+ Na + 1 - Cl 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 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 • • • K 1+ Mg 2+ Cu 2+ Cr 6+ Ba 2+ Hg 2+

Naming Anions • Anions are always the same. • Change the element ending to -- ide • F 1 - Fluorine to Fluoride

Name These • • • Cl 1 N 3 Br 1 O 2 I 1 Sr 2+ • • • Chloride Nitride Bromide Oxide Iodide Strontium

Write These • • Sulfide ion Iodide ion Phosphide ion Strontium ion • • S 2 I 1 P 3 Sr 2+

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). • SO 32 - = sulfite; SO 42 - = sulfate • Exceptions: Ammonium cation NH 4+, 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 • Na. Cl = Na+ Cl- = sodium chloride • Mg. Br 2 = Mg 2+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 Na 3 N Cr. N Sc. P Pb. O 2 Na 2 Se – – – – 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 – – Na. NO 3 Ca. SO 4 Cu. SO 3 (NH 4)2 O • Sodium nitrate • Calcium sulfate • Copper (II) sulfite • Ammonium oxide – – Li. CN Fe(OH)3 (NH 4)2 CO 3 Ni. PO 4 • 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. H Covalent bonds H N H + Cl Ionic bond - H 35

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 Ca 2+ – Chloride is Cl 1– Ca+2 Cl-1 would have a +1 charge – Need another Cl 1– Ca+2 Cl 2 -1 = Ca. Cl 2

Writing Formulas • Crisscross method Calcium chloride 2+ Ca 1 Cl Ca. Cl 2 No need to write the one Iron (III) sulfide Fe Fe 2 S 3 3+ 2 S Fe 2 S 3

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 • • • Li 2 S Sn. O 2 Mg. F 2 Cu. SO 4 Fe. P Fe 2 S 3 (NH 4)Cl (NH 4)2 S

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 Carbon Dioxide Water H 2 O CO 2 Ammonia NH 3

In a multiple covalent bond, more than one pair of electrons are shared between two atoms. (5. 2 e) • 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 N N (3 shared pairs) because nitrogen needs 3 electrons to fill its valence shell • Carbon dioxide has two double bonds 44

Regents Question: 08/02 #17 Which molecule contains a triple covalent bond? (1) H 2 (2) N 2 (3) O 2 (4) Cl 2 45

Molecular polarity can be determined by the shape of the molecule and the distribution of charge. • Possible shapes – Linear – Bent – Pyramidal – Tetrahedral (X 2 HX CO 2) (H 2 O) (NH 3) (CH 4 CCl 4) A polar molecule is called a dipole. It has a positive side and a negative side – uneven charge distribution. 46

Symmetrical (nonpolar) molecules include CO 2 , CH 4 , and diatomic elements. . . Symmetrical molecules are not dipoles. 47

Asymmetrical (polar) molecules include HCl, NH 3 , and H 2 O. (5. 2 l) The negative side of the molecule is the side that has the atom with the higher electronegativity. 48

Differences between ionic and covalent bonding: Ionic bonding • electron is “stolen” • high electronegativity difference • between metal & nonmetal • Formation of crystal structure think proportions of atoms in formula unit Na. Cl 1: 1 Na + Cl + Na + Cl -

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 2 3 4 5 6 7 8 monoditritetrapentahexaheptaocta- • 9 nona • 10 deca-

Naming Continued • To write the name…write two words 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 • • N 2 O NO 2 Cl 2 O 7 CBr 4 CO 2 Ba. Cl 2 H 2 O • • 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 • • P 2 O 5 I 4 O SF 6 NO 3 CH 4 PFl 3 Al. Cl 3

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 2. The number of electrons before you combine the atoms will equal number you have after.

Comparison of Bonding Types ionic covalent ions molecules nonconductive molten salts conductive transfer of electrons high mp DEN > 1. 7 valence electrons sharing of electrons low mp DEN < 1. 7

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. Strongest • • Ionic Solids Molecules with Hydrogen bonds Polar molecules Nonpolar molecules Weakest For nonpolar molecules, the greater the mass, the greater the force of attraction. 61

Hydrogen Bonds • Hydrogen bonds are considered to be dipole type interactions • Hydrogen bonds vary from about 4 k. J/mol to 25 k. J/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. CO 2 – 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 forces to be significant • Dipole-dipole forces are characteristically weaker than iondipole 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