University:
High School
Course:
High School Chemistry
Academic year:

2022
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492

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75
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larbiwxcvbn

Unit 4: Chemical Bonding & Molecules Chemical Bonding Chemical Bonds  Attraction between the nuclei and valence electrons of different atoms that ―glues‖ the atoms together.  the difference between materials as diverse as diamonds and pencils is how they're glued together.  Why?  Bonded atoms are more stable than solo atoms  How?  Atoms will share or exchange valence electrons to achieve a full outer shell (usually octet). 3 Main Types of Bonds Ionic Bonds – Transfer of electrons between atoms • • electrical attraction between cations & anions Formed by: metals & non-metals Covalent Bonds – sharing of electrons between atoms • • ―co‖ = sharing, ―valent‖ = outer electrons Formed by: non-metals & non-metals Metallic Bonds – Metal atoms that share a ―sea of electrons‖ • Formed by: metals & metals + - Predicting Bond Types  Bonding is not usually purely ionic or covalent, but somewhere in between  The difference in electronegativity strength of the atoms in a bond can help us estimate what percentage of the bond will be ionic (see example on next slide) Using the Periodic Table to Determine Bond Types Electronegativity Ionic bond= • metal (weak) & non-metal (strong) • huge difference in strength (1.7 or more) Metallic Bond = • 2 Metals (both weak) Covalent bond = • 2 non-metals (strong) •close to same strength Summary: Ionic Bonds vs. Covalent Bonds Ionic Covalent Electrons are Transferred (become charged ions that are attracted) Metal + non-metal (ex: Li + K ) Electrons are shared One atom is a lot higher electronegativity than the other (1.7) Close to equal electronegativities (less than 1.7) 2 non-metals ( O + O or O + N) Lewis Dot Structures Octet Rule – Most* atom wants to have 8 electrons in their valence shell (outermost shell) – Chemical compounds form so that each atom can complete their octet by gaining, losing or sharing electrons – *Exceptions = • H & He (they only want 2 electrons in their valence shell) • B (forms bonds so it will have only 6 electrons) • F, O & Cl (will sometimes be surrounded by more than 8 electrons because they are so electronegative) Lewis Dot Structure • Picture showing how many valence electrons an atom has (dots). •Helps determine how atoms will bond. Ex: Phosphorus (has 5 valence electrons) Lewis Dot Structures for Ionic Compounds • A way to show how atoms achieve the octet with each other. • Note: – the transfer of the electron – the charges ions that result This is how we draw it Lewis Dot Structures for Covalent Molecules 2 ways to show: – With electrons being shared in between Or H O H – Line showing the sharing of pair of electrons 3 Bonds Types in More Depth Covalent Bonds  Result from the sharing of electron pairs between two atoms  Molecule = termed used to describe atoms are held together by covalent bonds Covalent Bonds Occurs between 2 non-metals  electrons are shared 2 types of covalent bonds: Polar and non-polar (to be discussed later) • Ex: Water & most biological molecules (sugars, fats, proteins) •Can form single, double, or triple bonds Ionic Bonds •Forms between: Metal + Non-Metal •Electrons are transferred Ionic Bonds (cont.) Ex: Electroneg= .8 Electroneg= 3.0 Cl is so much stronger that it will “take” K’s electron The transfer of electron causes K to be a cation (+) and Cl to be an anion (-). Oppositely charged particles are highly attracted to each other… Ionic bond! Characteristics of Ionic Compounds Shape- crystal lattice of alternating positive and negative ions. Ex: NaCl and salts Ionic bonds are strong so they are:  hard  have a high melting point  high boiling point Crystal Lattice Metallic Bonds- ―sea of electrons‖ • Forms between 2 metals • Metal atoms valence electrons overlap creating a ―sea of electrons‖. •Electrons do not belong to any one atom, but roam freely throughout the metal atoms •Ex: Brass (alloy of Cu + Zn) Metallic Characteristics Because of these roaming ―sea‖ of electrons:  metals are great conductors of heat/ electricity  they are ductile (can be made into wire)  they are malleable (can be hammered into sheets) Electrical Conductivity Properties & Bonding Type Comparison Covalent Ionic Bonds Metallic Bonds Non-metal & non-metal Non-metal & metal Metal & metal Shared Transferred Sea of electrons Strong Very strong Varies Neutral group Crystal lattice crystalline Molecular Ionic metallic Melting Point Low Very high n/a Boiling Point Low High Very high Malleability n/a Not malleable, brittle Very malleable Ductility n/a Not ductile Very ductile Not conductive Conductive Highly conductive Formation Types of Atoms Electron Distribution Characteristics Bond Strength Structure Properties of Compounds Type of Compound Conductivity Bond Energy & Bond Length Bond Energy- energy required to break bond Bond Length Single Bond Bond Energy Low Double Bond Triple Bond High Bond Energy & Bond Lengths Bond H—Br H—C H—N H—O H—S C—O C=O C—C C=C CC O—O O=O N—N N=N NN Length (picometers) 141 109 101 96 93 143 129 154 134 120 148 121 145 125 110 Energy (kJ/mol) 366 413 391 464 339 360 799 348 614 839 145 498 170 418 945 Lewis Structures in Covalently Bonded Molecules & HONC Rule Drawing Lewis Dot Structures for Molecules arrange atoms to form a skeleton -Carbon is center atom -Hydrogen is never a central atom Pair up all electrons unpaired electrons can pair unpaired electron from another atom to form a bond Make sure each atom of the molecule obeys the octet rule & HONC rule Make sure you have correct # of valence electrons Examples of Lewis Dot Structure CH4 NH3 H H C H H H N H H H2O H O I2 I H I Multiple Covalent Bonds: Double bonds Two pairs of shared electrons O2 : CO2 : Multiple Covalent Bonds: Triple bonds Three pairs of shared electrons Molecular vs. Structural Formulas • Molecular formulas – show how many atoms of each element are in the molecules – Ex: C6H12O6 = 6 carbons, 12 hydrogens & 6 oxygens • Structural formulas – show the 2dimensional shape of the molecule – Ex: HONC 1-2-3-4 Rule • Hydrogen, oxygen, nitrogen & carbon are common elements found in biological molecules. – – – – Hydrogen needs 1 electron to fill its ―octet‖ Oxygen needs 2 electrons to fill its octet Nitrogen needs 3 electrons to fill its octet Carbon needs 4 electrons to fill its octet • ―1-2-3-4‖ can be used to predict how these atoms will form bonds with other atoms to build molecules. Molecular Geometry Seeing Molecules in 3-D Molecular Geometry molecules are really 3-D! CH4 in 2-D on a sheet of paper CH4 looks like this in 3-D Valence Electrons determine Molecular ―VSEPR‖ Shape • VSEPR = ―Valence-Shell Electron Pair Repulsion‖ • Electron pairs (bonding or lone pairs) in a molecule repel each other and will try and get as far away from each other as possible… this determines the shape. Lone pair electrons bonding pair electrons NH3 in 2-D NH3 VSEPR shape in 3-D 4 Shapes to Know Tetrahedral Pyramidal Bent Linear How Lone Pairs Affect Molecular Shape “paddles” are lone pairs of electrons. Remove the paddles and you can see the shapes. Steps for Determining Molecular Geometry 1. Draw Lewis dot structure 2. Count number atoms bonded to the central atom 3. Count number of lone-pair electrons on the central atom 4. Look up the Geometry on the chart Shapes in Large Molecules Large molecules are composed of the small shapes we’ve studied Ex: tetrahedral Why Shape Matters Ethyl Acetate (C4H8O2) Butyric Acid (C4H8O2) Same formula, but different shapes = very different smells Rum extract smell Rancid butter smell Polarity Differences In Electronegativities 3.3 Ionic 1.7 100 % 50% Polar-Covalent 0.3 0 5% Nonpolar-Covalent 0% Practice Problems Bonding Between: Difference in Electronegativity Bond Type Cl & Ca 3.0 – 1.0 = 2.0 O&H 3.5 – 2.1 = 1.4 Polar-covalent B&H |2.0 – 2.1| = 0.1 Nonpolarcovalent Ionic 2 types of Covalent Bonds: Non- Polar Polar (Arrow shows F is “pulling” electrons) -Electrons are shared equally - Usually the same element bonded to itself “partial positive charge” “partial negative charge” • Unequal sharing of electrons between atoms • more electronegative atoms ―hogs‖ electrons Visual Comparison of Bond Types Determining Polarity 1. Draw correct VSEPR Shape 2. Determine if molecule is symmetrical. 3. If the molecule is symmetrical = non-polar - no partial charges are needed! 4. If the molecule is NOT symmetrical = polar - you must show partial charges. - always bent or pyramidal shapes Ex: CO2 •Carbon dioxide= nonpolar • has polar bonds, but they cancel each other out. EX: Water= Polar Molecule How we know: 1) Cut the molecule on 2 planes - see how it’s different above the horizontal line = non symmetrical 1) One atom is ―pulling‖, look at periodic table to determine which one. + + Indicates which atom ―pulls‖ the electrons Means oxygen is slightly negative because it ―hogs‖ electrons 2 views of Polar water Non Polar Molecules • Non Polar molecule= ―no pull‖ – equal sharing of electrons – No difference in electronegativity – symmetrical in shape Cl Cl Examples of Polar & Nonpolar Molecules Inter vs. Intra molecular Forces Why Polarity Matters: Molecular Attractions • 1 molecule can be attracted to another molecule – ―inter‖molecular force • You can predict how one molecule might react with another: Ex: HBr + H2O Intermolecular Attractions & Smell • Besides shape, polarity also plays a role in your ability to smell. – Polar molecules = smell – Non-polar = don’t smell • Your smell receptors are polar and surrounded by mucous (a watery substance) Ex: Methane gas is odorless -They add a this stinky chemical to it so that you can smell it it: Intermolecular Forces vs. Intramolecular Forces Intramolecular Forces: (within in a molecule) Ex: -Covalent bond - Ionic bond -Metallic bond Intermolecular Forces: (between molecules) Ex: Hydrogen bonds -Weaker than covalent, ionic, & metallic bonds Hydrogen Bonding(an intermolecular force) in Water • Water is polar • (has a + and – end) •It’s ―sticky‖ • Will stick to any other thing that is: • polar (ex: other water molecules) • charged ionic substances (NaCl) Hydrogen bond Covalent bond Water’s Polarity leads to its ability to dissolve things so well The slight charges on water attract the NaCl’s ions and cause them to separate from each other Unique Properties of Water due to polarity & hydrogen bonding 1) Surface tension (hydrogen bonds create surface on water) 3) Adhesion/ Cohesion (water is attracted to other water molecules) 4) Capillary action water is attracted to other water molecules and will “rise” Properties of Water due to Hydrogen Bonding & Polarity • Cohesion – water molecules are attracted to one another – Causes water to be ―Sticky‖ – This is why water forms droplets • Adhesion – water is attracted to other substances – Water will ―stick‖ to containers & objects • Surface tension – strong forces between molecules cause the surface of a liquid to contract More properties… • Capillary Action – the movement of water within the spaces of a porous material due to the forces of adhesion, cohesion, and surface tension. •Universal Solvent things dissolve in water- polarity) Hydrogen Bonding in Kevlar Hydrogen bonding in Kevlar, a strong polymer used in bullet-proof vests. Hydrogen Bonding in DNA Other Intermolecular Forces (FYI… not part of this class) • Van der Waals Forces include: – Dipole-Dipole forces – results from the tendency of polar molecules to align themselves so that the positive end of one molecule is near the negative end of another molecule. – London (Dispersion) forces –results from the small, instantaneous dipoles that occur because of the varying positions of the electrons during their motion about nuclei Organic Chemistry Organic Chemistry- shows the versatility of carbon •has 4 valence electrons = 4 bonding spaces available. •Backbone to many large, complex biological molecules (Carbs, Lipids, Proteins, Nucleic Acids) •Over 16 million carboncontaining compounds are known. Monomers combine to make Polymers (small unit) (large) Common Examples of Polymers: Carbohydrates Lipids Proteins Nucleic Acids (CLPN) Ex: Carbohydrates Monomer Monosaccharide Polymer Polysaccharide Examples - Starch - Fiber - sucrose Ex: Lipids Monomers Glycerol & Fatty Acid tails Polymer Tri-glyceride Examples -Saturated Fats -Unsaturated fats -Steroids -Cholesterol Ex: Proteins Monomer Polymer Examples Amino Acids Polypeptide -enzymes -pigments -Meat/dairy Ex: Nucleic Acids Monomer Nucleotide Polymer Polynucleotide Examples -DNA -RNA Distilled Water vs. Tap Water Water Poisoning/ water Intoxication Cause: excessive consumption of water during a short period of time. Why: leads to a disruption in normal brain function due to the imbalance of electrolytes in the body’s fluids. – can dilute the careful balance of sodium compounds in the body fluids Who: individuals in water drinking contests…consume more than 10 liters (10.5 quarts) of water over the course of just a few minutes – People doing endurance sports which electrolytes are not properly replenished, yet massive amounts of fluid are still consumed Neural transmission