Chapter 5-The Structure of Matter Molecules and Compounds Ionic and Covalent Bonding Empirical and Molecular Formula

Compounds and Molecules

Remember from Chapter 2… A compound is a substance made of atoms of two or more different elements that are chemically combined by chemical bonds Mixtures are made of different substance that are just placed together; each substance in the mixture keeps its own properties

What is a chemical bond? An attractive force between the nuclei and valence electrons of different atoms that binds them together Bonds form in order to… decrease potential energy (PE) increase stability

Ne Stability Stability is the driving force behind bond formation! Octet Rule most atoms form bonds in order to have 8 valence e- full outer energy level like the Noble Gases! Ne Stability is the driving force behind bond formation!

Stability Transferring e- Sharing e-

H2O Chemical Formulas 1 oxygen atom 2 hydrogen atoms A compound always has the same chemical formula Shows: Elements in the compound The ratios of elements in the compound Example: H2O 1 oxygen atom 2 hydrogen atoms

Chemical Structure This tells you how the atoms or ions are arranged Described by bond length and bond angles Bond length: distance between the nuclei of two bonded atoms Bond angles: tell you how far away the atoms are away from each other in relation to the center atom Bond angles try to be as large as possible to decrease stress CO2 has bond angles of 180 degrees BF3 has bond angles of 120 degrees CH4 has bond angles of 109.5 degrees

Models of Compounds Space-filling: Ball and Stick: Shows the relative space occupied by oxygen and hydrogen Ball and Stick: can give you an idea of the geometry of a compound

How does structure affect properties? Some compounds exist as a large network of bonded atoms like quartz Rigid, strong structures Have high boiling and melting points Some compounds exist as a large network of bonded positive and negative ions Crystals are cubed shaped In NaCl, made of repeated network connected by strong bonds Made of tightly packed + and – ions Strong attractions cause high melting and boiling points Some compounds are made of many separate molecules Less strongly attracted atoms The strength of these attraction varies greatly

NaCl CO2 Vocabulary IONIC COVALENT Formula Unit Molecular Formula CHEMICAL FORMULA IONIC COVALENT Formula Unit Molecular Formula NaCl CO2

NaCl NaNO3 A. Vocabulary more than 2 elements 2 elements Binary COMPOUND more than 2 elements 2 elements Binary Compound Ternary Compound NaCl NaNO3

Na+ NO3- A. Vocabulary 1 atom 2 or more atoms Monatomic Ion Polyatomic

Types of Bonds

Ionic Bonds Ionic Bonds are formed by the TRANSFER of electrons What constitutes an ion? Positive charge=cation Negative charge=anion Opposites attract Cations will attract anions Anions will attract cations Bonding ALWAYS takes place in the valence shell between valence electrons

Properties of ionic bonds Strong bonding forces Hard, brittle Crystalline structure Conduct electricity in solution High melting/boiling point

What makes an ionic bond? Must include a metal + a non-metal Metal = cation Non metal = anion Anion “takes” electron from cation Cation is willing to give it away, to lower an energy level Anion is looking to fulfill the octet rule For this reason – group I and II metals “LOVE” group 17 non-metals (halogens) Creates a salt

Covalent Bonds Compounds that are made of molecules Bonds that are formed by SHARING electrons Electrons are shared from the valence shell, still want it full! But electrons are not always shared equally… When two Chlorine atoms bond, electrons are equally attracted to the positive nucleus of each atom---NON- POLAR Covalent bond When not shared equally-Polar Covalent bond

Properties of Covalent Bonds Strong bonding forces Can be found as liquids, gases, or soilds Do not conduct electricity Low melting/boiling point Non-metal + non-metal All the polyatomics are covalently bonded

What makes a Covalent Bond Non-metal + non-metal Some examples below Can share more than one pair of electrons single bond shares one pair electrons double bond shares two pairs electrons triple bond shares three pairs electrons

Naming Covalent Compounds Write the more metallic element first. Add subscripts according to prefixes.

Covalent Bonds Two nonmetal atoms form a covalent bond because they have less energy after they bonded H + H H : H = HH = H2 hydrogen molecule

Diatomic Molecules Gases that exist as diatomic molecules Are: Br2, I2 , N2, Cl2, H2, O2, F2 octets          N  +  N   N:::N   triple bond

Molecular Formulas The Seven Diatomic Elements Br2 I2 N2 Cl2 H2 O2 F2

Comparison Chart IONIC COVALENT Bond Formation The “glue” of the bond e- are shared between two nonmetals e- are transferred from metal to nonmetal The “glue” of the bond electrical attraction Physical State solid liquid , gas, or solid Melting Point high low Solubility in Water yes usually not Electrical Conductivity yes no Other Properties crystal lattice shape = brittle; exothermic odor

Metallic Bonds Metal + metal Attraction between one atom’s nucleus and a neighboring atom’s electrons Closely packed structure, causes outermost energy level (valence shell) to overlap Electrons are free to move from atom to atom This is why they are good conductors of electricity

Metallic Bond Chart METALLIC “electron sea” solid very high no e- are delocalized among metal atoms Bond Formation “electron sea” Type of Structure Physical State solid Melting Point very high Solubility in Water no Electrical Conductivity yes (any form) Other Properties malleable, ductile, lustrous

Hydrogen Bonds Explains why an oxygen of one water molecule is attracted to hydrogen atom of a neighboring water molecule Water molecules attract each other, but these bonds aren’t as strong as the bonds holding oxygen and hydrogen molecules together within a molecule

Writing Lewis Dot Structures for Covalent Compounds

Lewis dot structres Na Cl = Simple depiction to help us “visualize” the location of the valence electrons Only includes valence electrons Includes element symbol and dots representing valence electrons Na Cl = Loses electron; becomes + Gains electron; becomes -

What’s the charge? The charge on an ion. Indicates the # of e- gained/lost to become stable. 1+ 4- 2+ 3+ 4+ 3- 2- 1-

How do ions bond? +1 -1 Review: cations + anions What is the ionic charge of Na? +1 What is the ionic charge of Cl? -1 The final compound MUST be neutral!!

NASL Method STEPS Figure out what element is central (I will tell you this part) Calculate the N (needed) number of electrons for all atoms to abide by the octet rule (except 2 for H & He) Calculate the A (available) number of electrons by adding up all the valence electrons for each atom Calculate the S (shared) number of electrons = N-A Calculate lone L (lone pairs or dots) as the difference between A – S. Check to make sure you have used as many electrons in A step

H2O N A S L

SO2 N A S L

SOCl2 N A S L

O2 N A S L

What if it’s charged? STEPS When Calculating N include charge Anions GAIN that many electrons Cations LOSE that many electrons Calculate A, S, L just the same When finished with lewis dot structures, draw brackets around picture with the charge on the outside

(NH4)+1 N A S L

(NO3)-1 N A S L

(SO4)-2 N A S L

Empirical and Molecular Formulas

C2H6 CH3 Empirical Formula Smallest whole number ratio of atoms in a compound C2H6 reduce subscripts CH3

Empirical Formula 1. Find mass (or %) of each element. 2. Find moles of each element. 3. Divide moles by the smallest # to find subscripts. 4. When necessary, multiply subscripts by 2, 3, or 4 to get whole #’s.

B. Empirical Formula Find the empirical formula for a sample of 25.9% N and 74.1% O. 25.9 g 1 mol 14.01 g = 1.85 mol N = 1 N 1.85 mol 74.1 g 1 mol 16.00 g = 4.63 mol O = 2.5 O

N2O5 N1O2.5 B. Empirical Formula Need to make the subscripts whole numbers  multiply by 2 N2O5

CH3 C2H6 C. Molecular Formula empirical formula molecular formula ? “True Formula” - the actual number of atoms in a compound Different compounds can have the same empirical formula CH3 empirical formula ? C2H6 molecular formula

C. Molecular Formula 1. Find the empirical formula. 2. Find the empirical formula mass. 3. Divide the molecular mass by the empirical mass. 4. Multiply each subscript by the answer from step 3.

(CH2)2  C2H4 C. Molecular Formula empirical mass = 14.03 g/mol The empirical formula for ethylene is CH2. Find the molecular formula if the molecular mass is 28.1 g/mol? empirical mass = 14.03 g/mol 28.1 g/mol 14.03 g/mol = 2.00 (CH2)2  C2H4