Life Update

So I have made my final decision about this part of my life...I am moving to Curacao. 

I will be leaving for the states June 25th. I'll be bringing my dog with me because my parents are going to watch her for me while I'm there. I will be there for 8 months, and then move back to the states for my internado. I will hopefully be able to come back to the DR to visit. 

I have been so honored to be your teacher this year and will miss you all so much. If you want to follow my future adventures, I have another blog with occasional updates about my life and where I am at. Here is the link to my other blog Kaitlyn's Adventures. Otherwise, feel free to email me if you miss me or have questions: kaitlyn.cunningham@icloud.com or nyltiak@gmail.com.

Thank you for everything this year. I will never forget this experience. 

Miss Cunningham

Answers to Practice

Good afternoon all,

Here are the answers to the practice. I also have a list of the equations that you are allowed to bring to the final. You can bring a the equations for the pH chapter and titration.

Let me know if you have any questions about the practice. Email me at kaitlyn.cunningham@icloud.com or nyltiak@gmail.com with any questions.

Practice for Final Exam

1. 9.35g of Al₂O₃ contains how many molecules?
2. 0.253 moles of BN has a mass of…
3. Calculate the number of moles of 8.46 x 10²⁴ atoms of F.
4. Calculate the number of moles of 4.35g of Cs₂CO₃.
5. 0.692moles of CsH contains how many molecules?
6. How many moles are in 14.49 g of CoSO₄.
7. What is the mass of 7.20 x 10²³ atoms of Ne?
8. 50.82 g of Cr₂O₃ has how many moles?
9. 2.84 moles of N₂H₄ has how many molecules?
10. 5.29 x 10²⁴ molecules of HCl has a mass of…
11. In a 30.0 g sample of Pb(NO₃)_2, Pb has a mass of 2.535 g, N a mass of 8.694 g, and O a mass of 18.768 g. What is the % composition?
12. In a 25.0 g sample of Hg₂Cl₂, Hg has a mass of 3.755 g and Cl has a mass of 21.245 g. What is the % composition.
13. In a 20.0 g sample of Ni(NO₃)₂, Ni has a mass of 3.066 g, Ni has a mass of 6.426 g, and O has a mass of 10.508 g. What is the % composition?
14. What is the % composition of the following:
• KMnO₄
• Sc₂O₃
• H₂SeO₄
• Cl₃HSi
• NaBrO
15. What volume of 0.075 M HCl is required to neutralize 100 mL of 0.01 M NaOH solution?
16. What is the concentration of NaOH if 45.2 mL of 1.5 M H₂SO₄ is needed to neutralize 20.0 mL of NaOH: H₂SO₄ + 2NaOH ➝ Na₂SO₄ + 2H₂O?
17. Find the [H₃O⁺] and pH of the following: 0.075 M HCl 
18. Find the [H₃O⁺] and pH of the following: 0.122 M H₂SO₄
19. Calculate the [H₃O⁺] and pOH of the following solution: pH = 9.5
20. Calculate the [H₃O⁺] and pOH of the following solution: [OH^-] = 3.9 x 10^-12
21. Find the [OH^-] of the following solutions: pOH = 1.5, pH = 4.6, & [H₃O⁺] = 2.5 x 10^-9
22. In the following chemical reactions, identify the acid, base, conjugate acid, and conjugate base.
• HC₂H₃O₂ + H₂O ⇌ H₃O⁺ + C₂H₃O₂^-
• NH₃ + H₂O ⇌ NH₄⁺ + OH^-
• HNO₃ + H₂O ⇌ NO₃^- + H₃O⁺
23. Find the [H₃O⁺] and pH of the following:
• 0.025 M HNO₃
• 3.4 x 10^-4 M H₂SO₄
24. Calculate the [H₃O⁺] and pOH of the following solutions:
• pH = 3.2
• pH = 5.0
• [OH^-] = 8.2 x 10^-9
• pH = 12.4
25. Find the [OH^-] of the following solutions:
• [H₃O⁺] = 3.9 x 10^-6 M
• [H₃O⁺] = 0.0014 M
• pH = 4.2
26. If 35.2 mL of 12M HCl is used to titrate 50.0 mL of KOH, what is the concentration of the KOH?
27. What is the volume of 2.5 M NaOH needed to titrate 25.0 mL of 1.0 M HCl?
28. Given the following equation, find the concentration of NaOH when 24.09 mL of 1.605 M H₂SO₄ is needed to titrate 50.0 mL of NaOH.      H₂SO₄ + 2NaOH ➝ Na₂SO₄ + 2H₂O
29. What is the salt that is formed during the following neutralization reactions?
• Mg(OH)₂ + HCl ➝ ? + H₂O
• H₂SO₄ + 2NH₄O ➝ ? + 2H₂O
• Ni(OH)₂ + 2HClO₄ ➝ ? + 2H₂O
• Mg(OH)₂ + H₂SO₄ ➝ ? + 2H₂O
30. Balance the following redox reactions:
• Na + H₂O → NaOH + H₂
• HCl + HNO₃ → HOCl + NO + H₂O (*Cl on the product side has an oxidation number of +1)
• SnCl₄ + Fe → SnCl₂ + FeCl₃
• CO + I₂O₅ → I₂ + CO₂
• MnO₄^- + Fe²⁺ + H⁺ → Fe³⁺ + Mn²⁺ + H₂O
• Cu⁺ + Fe → Fe³⁺ + Cu

Proteins and Nucleic Acids

Proteins

• Most important substance in the bodies of living things.
• Building blocks for muscle, hair, blood cells, skin, slk, enzymes, insulin, etc.

Amino Acids

• Building blocks of proteins.
• Contain an amine (NH2) group and carboxyl (COOH) group.
• Various side chains (R groups) attach to the carbon adjacent to the N.
• Results in 20 different common amino acids.
• Our bodies can synthesize 12 out of 20.
• The other 8 must come from diet
• Known as essential amino acids

List of Amino Acids

Polypeptide Chains

• Amino acids joined together with peptide bonds.
• Dipeptides: two amino acids joined together
• Aspartame (artificial sweetener)
• Polypeptides: polymers of many amino acids

Polypeptide chain or protein

Proteins

• Polymer of one or more polypeptide chains.
• Some contain several hundred amino acids others several thousand.
• Various structures

Nucleic Acids

DNA Structure

• When cells reproduce, they pass genetic information to one another in long-chain molecules called chromosomes.
• Human body contains 46 chromosomes in the nuclei.
• Segments of chromosomes, called genes, carry the coded information to that directs the production of specific polypeptide chains.
• Nucleotides are the building blocks of nucleic acid.
• Three units:
• 5 carbon sugar (ribose or dioxyribose)
• phosphate group
• ring-shaped nitrogen containing base
• There are five possible nitrogen-congaing bases attached to the sugar:
• Adenine (A)
• Cytosine (C)
• Guanine (G)
• Thymine (T)
• Uracil (U)
• DNA has A, C, G, T
• RNA has A, C, G, U
• DNA is two strains coiled around each other in a double helix
• When cells divide, a DNA strand reproduces by first unraveling with the help of enzymes. 
• Each of the two resulting strands serves as a template, or pattern, for a new complementary chain.
• Complementary bases on the newly forming strand match the now-exposed bases of the old strand.
• When the process of replication is completed, two identical double-stranded DNA molecules result. 

DNA Replication

• One strand goes to each half of the dividing cell.
• DNA and RNA molecules instruct cells on how to make proteins.
• DNA is used as a template to make RNA.
• RNA is made up of nucleotides.
• Trinucleotide groups (3 nucleotides) code for various amino acids.
• As the amino acids join together, they form proteins. 
• The proteins are the chemical manifestation of our physical characteristics. 

This is an example of protein synthesis from RNA

Differences between DNA and RNA

• DNA
• the sugar is deoxyribose
• contains thymine
• double strand
• RNA
• the sugar is ribose
• contains uracil
• single strand

Carbohydrates

• Most abundant biological compounds.
• Sugars and starches make up a large part of the human diet.
• Each year photosynthetic processes in plants convert water and carbon dioxide into one hundred billion tons of carbohydrates.
• Exoskeletons of insects are made of carbohydrates.

• All carbohydrates have several of the –OH (hydroxy) groups common to alcohols.
• They also have the C=O (carbonyl) group of aldehydes and ketones.
• They are known as one or more polyhydroxy aldehyde or ketone
• Three main functions:
• energy storage
• an energy source for cellular functions
• glucose
• structural elements in plants and animals
• cellulose & chitin
• Classified into three groups based on the number of sugar units they contain.
• Monosaccharides 
• Disaccharides
• Polysaccharides

Monosaccharides

• These are simple sugars. One polyhydroxy aldhyde or ketone.
• Rarely occur in nature by themselves.
• Usually bonded to an protein, fat or other carbohydrate.
• Glucose is the most abundant sugar in nature.
• The human body maintains a reasonably constant level of 80 to 120 mg of glu- cose per 100 mL of blood.
• Glucose is also the fundamental building block of the most common long-chain carbohydrates.
• Glucose in an aqueous solution exists in an equilibrium between the ring form and the straight-chain form.
• Fructose also forms ring structures when it is in solution.
• Since the carbonyl group is not at the end of the carbon chain, it forms a five-member ring.

Disaccharides

• Contain two monosaccharide units.
• An oxygen bridge between the two monosaccharides holds the two units together.
• Three disaccharides play an important part in the human diet—maltose, lactose, and sucrose.
• Maltose (glucose - glucose)
• germinating grain & digestion of starches
• Lactose (glucose - galactose)
• milk
• Sucrose (glucose - fructose)
• table sugar, fruits, sugar cane, beets, nectar

Polysaccharides

• Many sugar units, up to several million units
• Can be built with several different monosaccharide units
• Glycogen: energy storage 
• Cellulose: gives plants structure
• Starch: supply nearly 3/4ths of the worlds food energy

Quiz on Wednesday

Study:

• Difference between aliphatic and aromatic
• Difference between alkanes, alkenes, alkynes
• What is a organic compound
• What is unique about carbon
• Nomenclature

Hydrocarbons

What are hydrocarbons?

• The simplest aliphatic and aromatic compounds
• Only contain hydrogen and carbon
• Can contain single, double, or triple bonds
• Classified by the types of bonds.

Alkanes

• Alkanes are aliphatic hydrocarbons that only contain single bonds.
• The general formula for alkanes is C_nH_2n+2
• Example: CH₄, C₂H₆, C₃H₈
• Considered saturated because carbon is surrounded by the max number of hydrogen.
• Methane is the simples alkane (CH4)
• Each consecutive alkane adds a carbon and its respective hydrogens.
• Methane: 1 carbon
• Ethane: 2 carbons
• Propane: 3 carbon
• Butane: 4 carbons
• Properties of Alkanes
• Very low melting and boiling points. They rise as carbons are added.
• Ex. methane
• Melting point (C):  -183
• Boiling point (C): -164
• Non-polar

Alkenes

• Hydrocarbons that contain double bonds between carbon atoms
• Contain the prefix -ene.
• The smallest is ethene. Why not methene?
• Considered unsaturated, because the double bond prevents the max number of hydrogen from bonding.
• Naming of alkenes requires numbering the carbons to identify the place where the double bond is.
• You start at the carbon that will give you the smallest number.
• This allows us to know where the double bond is. 
• Properties of Alkenes:
• Slightly higher melting/boiling points
• The first couple are gases at room temperature
• Relatively non polar

Alkynes

• Hydrocarbons that contain triple bonds between the carbon atoms
• Uses the prefix -yne
• The simplest is the most common, ethyne (acetylene).

Cyclic Aliphatic Compounds

• Not all hydrocarbons are open chains of carbon atoms.
• Some form a ring.
• 5 and 6 alkane rings are most abundant. 
• Some can have more than one double bond.

Aromatic Structures

• All contain a form of a molecule benzene
• They are called aromatic because they often smell good
• C₆H₆ is the simplest aromatic compound known
• It was hard to figure out the structure:
• Behaves like an alkane, but they knew from the molecular weight that it had several double and triple bonds.
• When they measured the bond length, the found that it should contain 1.5 bond lengths.
• Showed that carbon was in a ring and all had identical bonds
• In 1865, August Kekule proposed the structure. 
• He said that it was a dynamic equilibrium of the two.
• The double bonds were not “tied dow”, but are more or less shared.

Organic Chemistry Introduction

• Organic Compounds: covalently bonded carbon compounds, with the exception of carbonates, carbon oxides, and carbides.
• Biochemistry: the study of complex reactions taking place between organic compounds within living organisms.

Unique Carbon Atom

• Carbon has some unique properties that enable it to form hundreds of thousands of compounds.
• Carbon has 4 valence electrons, requiring 4 bonds to obtain an octet
• Carbon forms strong chemical bonds with other carbon atoms
• Carbon forms stable, almost non polar bonds with hydrogen
• Carbon atoms can bond to a wide variety of atoms
• H, P, O, N, S, the halogens, and even metal atoms.
• Bonds can be straight, branched, and in various lengths.
• They can even form rings
• Can form double and triple bonds

Structural Forumlas

• Structural formulas are used a lot in organic chemistry because molecular formulas can mean various compounds. 
• C₂H₆O can mean ethanol or dimethyl ether
• See page 446 in your books

Classification

• There are approx. 300,000 new organic compounds synthesized for the first time every year.
• It is important to have some categories:
• Aliphatic compounds: without a benzene ring
• Aromatic compounds: with a benzene ring

Redox Reactions

• Oxidation-reduction reactions (redox reactions):
• Reactions involving transfers or shifts of electrons 
• Oxidation:
• A loss of electrons which makes the oxidation number go up
• Occurs mainly in metals
• Occurs in some covalently bonded substances
• Does not require oxygen (that is not what oxidation means!)
• Reduction:
• A gain of electrons that makes the oxidation number go down (reduced)
• Occurs mainly in nonmetals that gain electrons by taking them from metals
• Review of oxidation numbers:
• Rule 1: free atoms = 0
• Rule 2: ion charge = oxidation number
• Rule 3: compound sum = 0
• Rule 4A: Group 1 = +1
• Rule 4B: Group 2 = +2
• Rule 4C: H = +1 or -1
• Rule 4D: O = -2 or -1
• Rule 4E: Group 17 = -1
• Rule 5: sum of ONs in polyatomic ion = charge
• Practice: Assign Oxidation Numbers
• H₂CO₃
• H: +1, O: -2, C: +4

• N2
• N: 0
• Zn(OH)₄^-2
• Zn: +2, H: +1, O: -2

Redox

• Short for reduction­-oxidation
• Pronounced “REE-docs”
• Must occur together (An element cannot take electrons without another one losing them.)
• LEO the GERm
• Lose Electrons Oxidation
• Gain Electrons Reduction
• Determine which element is oxidized and which is reduced?
• Zn + 2H⁺ ➝ Zn²⁺ + H₂
• Zn is oxidized (ON: 0 ➝ +2)
• H⁺ is reduced (ON: +1 ➝ 0)
• 3Hg²⁺ + 2 Fe(s) ➝ 3Hg + 2Fe³⁺
• Hg²⁺: reduced 
• Fe: oxidized
• Oxidizing and Reducing Agents:
• Reducing agent is a substance used to reduce another substance.
• If a substance is oxidized it is the reducing agent.
• Oxidizing agent is a substance used to oxidize another substance.
• If a substance is reduced it is the oxidizing agent.
• Example:
• Zn + 2H⁺ ➝ Zn²⁺ + H₂
• Zn is oxidized (ON: 0 ➝ +2) [REDUCING AGENT]
• H⁺ is reduced (ON: +1 ➝ 0) [OXIDIZING AGENT]
• 3Hg²⁺ + 2 Fe(s) ➝ 3Hg + 2Fe³⁺
• Hg²⁺: reduced [OXIDIZING AGENT]
• Fe: oxidized [REDUCING AGENT]

Balancing Redox Reactions

See practice for the procedure, but here are the steps:

1. Assign oxidation numbers
2. Make half reactions (oxidized reaction and reduced reaction) [forget everything else for now]
3. Balance electrons
4. Add everything back in and balance traditionally [see Chapter 8 for procedure]

Answers to Chapter 17 Practice 28/03

Here are the answers and more or less how to do the process. If you have any questions, you can email me or ask on Monday.

Oxidized/Reducing Agent: Na 

Reduced/Oxidizing Agent: H₂O 

Oxidized/Reducing Agent: HCl 

Reduced/Oxidizing Agent: HNO₃

Oxidized/Reducing Agent: Fe

Reduced/Oxidizing Agent: SnCl₄

Oxidized/Reducing Agent: CO

Reduced/Oxidizing Agent: I₂O₅

Oxidized/Reducing Agent:  Fe²⁺

Reduced/Oxidizing Agent: MnO₄^-

Oxidized/Reducing Agent: Fe

Reduced/Oxidizing Agent: Cu⁺

Review for Test

1. Definitions of acids/bases.
• Properties
• Arrhenius defnition
• Bronsted-Lowry definiton
• Conjugate pairs
• Lewis definition
2. Acid Base Equilibria
3. Self-Ionization of Water (K_w)
4. pH, pOH, [H₃O⁺], [OH^-]
5. pH scale and values
6. Acid-Base Strength (Concentration vs. Strength)
• Ka and Kb
7. Polyprotic Acids
8. Titration
9. Buffers
10. Important Equations:
• Kw = [H₃O⁺][OH^-] = 1.0 x 10^-14
• pH = -log[H₃O⁺]
• [H₃O⁺] = 10^-pH
• pOH = -log[OH^-]
• [OH^-] = 10^-pOH
• pH + pOH = 14
• (M_K)(V_K)=(M_U)(V_U)

1. In the following chemical reactions, identify the acid, base, conjugate acid, and conjugate base.
• HC₂H₃O₂ + H₂O ⇌ H₃O⁺ + C₂H₃O₂^-
• NH₃ + H₂O ⇌ NH₄⁺ + OH^-
• HNO₃ + H₂O ⇌ NO₃^- + H₃O⁺
2. Find the [H₃O⁺] and pH of the following:
• 0.025 M HNO₃
• 3.4 x 10^-4 M H₂SO₄
3. Calculate the [H₃O⁺] and pOH of the following solutions:
• pH = 3.2
• pH = 5.0
• [OH^-] = 8.2 x 10^-9
• pH = 12.4
4. Find the [OH^-] of the following solutions:
• [H₃O⁺] = 3.9 x 10^-6 M
• [H₃O⁺] = 0.0014 M
• pH = 4.2
5. If 35.2 mL of 12M HCl is used to titrate 50.0 mL of KOH, what is the concentration of the KOH?
6. What is the volume of 2.5 M NaOH needed to titrate 25.0 mL of 1.0 M HCl?
7. Given the following equation, find the concentration of NaOH when 24.09 mL of 1.605 M H₂SO₄ is needed to titrate 50.0 mL of NaOH.      H₂SO₄ + 2NaOH ➝ Na₂SO₄ + 2H₂O
8. What is the salt that is formed during the following neutralization reactions?
• Mg(OH)₂ + HCl ➝ ? + H₂O
• H₂SO₄ + 2NH₄O ➝ ? + 2H₂O
• Ni(OH)₂ + 2HClO₄ ➝ ? + 2H₂O
• Mg(OH)₂ + H₂SO₄ ➝ ? + 2H₂O

Answers:

1. Identifying the acid, base, conjugate acid, and conjugate base.
• Acid: HC₂H₃O₂Base: H₂OC. Acid: H₃O⁺C. Base: C₂H₃O₂^-
• Acid: H₂OBase: NH₃C. Acid: OH^-C. Base: NH₄⁺ 
• Acid: HNO₃Base: H₂OC. Acid: H₃O⁺C. Base: NO₃^- 
2. Find the [H₃O⁺] and pH of the following:
• [H₃O⁺] = 0.025 & pH = 1.6
• [H₃O⁺] = 6.8 x 10^-4 & pH = 3.1
3. Calculate the [H₃O⁺] and pOH of the following solutions:
• [H₃O⁺] = 6.0 x 10^-4 & pOH = 10.8
• [H₃O⁺] = 1.0 x 10^-5 & pOH = 9
• [H₃O⁺] = 1.2 x 10^-6 & pOH = 8.1
• [H₃O⁺] = 4.0 x 10^-13 & pOH = 1.6
4. Find the [OH^-] of the following solutions:
• [OH^-] = 2.5 x 10^-9
• [OH^-] = 7.1 x 10^-12
• [OH^-] = 1.6 x 10^-10
5. 8.4 M KOH
6. 10.0 mL of NaOH
7. 1.5 M NaOH
8. What is the salt that is formed during the following neutralization reactions?
• MgCl₂
• (NH₄)₂SO₄
• Ni(ClO₄)₂
• MgSO₄

Acid Base Equilibrium

• Why can we eat and drink some acids and bases and not others?
• The degree to which they release or accept protons
• Equilibrium constants describe how readily acids deprotonate and bases protonate.

Self-Ionization of Water

Water can react with itself.

One molecule accepts and one donates a proton. 

This is called self-ionization

• This reaction is very important. Because it gives us a constant for acid base equilibrium. Kw is the equilibrium constant for water. 
• Kw = [OH^-][H₃O⁺] = 1.0 x 10^-14 M
• Note: M stands for molarity which is moles/L
• Whether a solution is acidic, basic or neutral, the product of the [H3O+] and [OH-] is always equal to Kw.
• Example problem: 
• The [H₃O⁺] in a mild acid is found to be 5 x 10^-7 mol/L. What is the concentration (molarity) of hydroxide ions?
• [OH^-][H₃O⁺] = 1.0 x 10^-14 M
• [OH^-] = 1.0 x 10^-14 M/[H₃O⁺]
• [OH^-] = 1.0 x 10^-14 M/5.0 x 10^-7 M
• [OH^-] = 2.0 x 10^-8 M

pH Scale

• pH stands for “power of hydronium”
• pH is the negative logarithm of the [H₃O⁺]
• pH = -log [H₃O⁺]
• Examples:
• If [H₃O⁺] = 0.0025 M
• pH = -log(0.0025) = 2.6
• If [H₃O⁺] = 4.57 x 10^-9 M
• pH = -log(4.57 x 10^-9 M) = 8.34

pH Values

• A pH of 7 means neutral.
• A pH of 0-7 is an acid.
• A pH of 7-14 is basic

pH Scale

• Example 1:
• The [H₃O⁺] in a shampoo is 2.0 x 10^-5 M. What is the pH of this shampoo?
• pH = -log [H₃O⁺]
• pH = -log(2.0 x 10^-5) = 4.7
• Example 2:
• What is the pH of an aqueous solution of 0.40 g of HI dissolved in 500 mL of water?
• First, convert grams of HI to moles of HI.
• 0.4 g HI x (1 mol HI/127.9 g HI) = 0.00031 mol HI
• Next calculate molarity (M).
• (0.0031 mol HI/500 mL) x (1000 mL/1L) = 6.3 x 10^-3 M 
• Solution:
• pH = -log[H₃O⁺
•       = -log(6.3 x 10^-3 M)
•       =  2.2
• Example 3
• Find the pH of a solution whose [H₃O⁺] equals 9.5 x 10^-8.
• pH = -log[H₃O⁺]
•       = -log(9.5 x 10^-8 M)
•       =  7.02 

Defining Acids and Bases

• We live in a world full of chemicals. Some of these chemicals are compounds called acids and bases. We will talk about the properties and various ways of identifying acids and bases. In 1663, Robert Boyle started to classify compounds as acids and bases by their physical properties, but it would be another 200 years before scientists began to explain their chemical properties.

Properties of Acids

• Sour taste:
• citric acid makes fruit juice tart
• acetic acid makes pickles sour
• acetylsalicylic acid makes aspirin sour
• lactic acid gives the smell and flavor to sour milk
• When some acids dissolve, they are able to conduct and electrical current. For example, HCl ionizes into hydrogen ions and chlorine ions. Any substance that ionizes to conduct electricity in a solution is called an electrolyte.
• An acid turns blue litmus paper red.
• Acids react with active metals to produce hydrogen gas and a salt.
• Mg (s) + 2HCl (aq) → MgCl₂ (aq) + H₂ (g)

Properties of Bases

• Bitter taste
• Slippery
• Red litmus paper turns blue.
• Neutralization reaction between an acid and a base in an aqueous solution produces a salt and water.
• HCl (aq) + NaOH (aq) → NaCl (aq) + H₂O

The Arrhenius Model

• Acids: 
• Give a H⁺ ion (a proton) when in water.
• This hydrogen ion is not stable so it bonds with the water to form the hydronium ion (H₃O⁺)
• Examples:
• HCl + H₂O → H₃O⁺ + Cl^-
• CH₃COOH + H₂O → H₃O⁺ + CH₃COO^-
• Bases:
• Gives an OH^- ion (a hydroxide) when in water
• Examples:
• NaOH + H₂O → OH^- + Na⁺ + H₂O
• NH₃ + H₂O → OH^- + NH₄⁺
• Limitations:
• Not all substances with Hs or OHs are an acid or a base.
• Examples: CH₄ and CH₃OH
• Only works for acids and bases in water.

The Bronsted Lowry Model

• Acids: donate proton(s)
• This is the same definition as in the Arrhenius model.
• Donating or losing a proton is called deprotonation.
• Bases: accept proton(s)
• Much wider than Arrhenius because it includes substances that might not have an OH group.
• Gaining or accepting a proton is called protonation.
• Conjugate pairs:
• This refers to acids and bases with common features. These common features are the equal loss/gain of protons between pairs. Conjugate acids and conjugate bases are characterized as the acids and bases that lose or gain protons.
• Acid + base → conjugate base + conjugate acid
• Example:
• HC₂H₃O₂ + H₂O <—> H₃O⁺ + C₂H₃O₂^-
• Acid: HC₂H₃O₂ because it donates a proton
• Base: H₂O because it accepts the proton
• Conjugate acid: H₃O⁺ because it would donate a proton
• Conjugate base: C₂H₃O₂^- because it would accept the proton
• HClO₂ + H₂O → ClO₂^- + H₃O⁺
• Acid: HClO₂
• Base: H₂O
• Conjugate acid: H₃O⁺
• Conjugate base: ClO₂^- 
• OCl^- + H₂O → HOCl + OH^-
• Acid: H₂O
• Base: OCl^- 
• Conjugate acid: HOCl
• Conjugate base: OH^-
• HCl^- + H₂PO₄^-  → Cl^- + H₃PO₄
• Acid: HCl^- 
• Base: H₂PO₄^-
• Conjugate acid: H₃PO₄
• Conjugate base: Cl^- 

Lewis Model

• Acids: accept a pair of e- (one empty orbital)
• Bases: donate a pair of e- (one unbonded pair of electrons)