Chemistry

Matter

• Organisms are composed of matter

• Matter is anything that takes up space and has mass

• Matter is composed of chemical elements

• Matter is found on the Earth in three physical states

– Solid

– Liquid

– Gas

States of Matter• Gases take the shape and volume of their container and can be compressed to

form liquids.

• Liquids take the shape of their container, but they do have their own volume

• Solids are rigid and have a definite shape and volume.

Classification of Matter• Element: a substance composed of only one type of atom (all the

atoms have the same number of protons).

• Molecule: a unit composed of two or more atoms joined together by chemical bonds

• Compound: a substance composed of 2 or more elements that have been joined by chemical bonds

• Mixture: a combination of 2 or more substances that do NOT chemically bond e.g. sugar mixed with salt

Sodium Chlorine Sodium chloride

Pure Substances and Mixtures

Elements• If a pure substance cannot be decomposed into something else, then the substance is an element

• There are 114 elements known, 92 naturally occurring

• Each element is given a unique chemical symbol (one or two letters)

Periodic Table

Essential Elements Of Life• Only about 25 of the elements are essential to life

• Carbon, hydrogen, oxygen, and nitrogen make up 96% of living matter

• Most of the remaining 4% consists of calcium, phosphorus, potassium, and sulfur

• Trace elements are those required by an organism in minute quantities

Essential Elements Of Life

Periodic Chart

Atoms• Each element consists of one kind of unique atom

• An atom is the smallest unit of matter that still retains the properties of an element, it cannot be broken down to other substances by chemical reactions

Subatomic Particles• Atoms are composed of subatomic

particles

• Relevant subatomic particles include:

– Neutrons (no electrical charge)

– Protons (positive charge)

– Electrons (negative charge)

• Neutrons and protons form the atomic nucleus

• Electrons form a cloud around the nucleus

Nucleus

Cloud of negativecharge (2 electrons)

(a)

(b)

2 Protons

2 Neutrons

2 Electrons

Atomic Number And Atomic Mass

• Atoms of the various elements differ in number of subatomic particles

• An element’s atomic number is the number of protons

• The number of protons (atomic number) determines the element’s properties

• An element’s mass number is the sum of protons plus neutrons in the nucleus

• Atomic mass, the atom’s total mass, can be approximated by the mass number

Atomic number

Element symbol

Mass number

Periodic Chart

Orbitals• Electrons orbit the nucleus of an atom in specific electron shells

• Each Orbital holds a maximum of 2 electrons each

• Several orbitals may be the same distance from the nucleus and thus contain electrons of the same energy. Such electrons are said to occupy the same energy level or shell.

• Rule of Eights for filling each shell:

Electron

First electron shell(can hold 2 electrons)

Outermost electron shell(can hold 8 electrons)

Carbon (C)Atomic number = 6

Nitrogen (N)Atomic number = 7

Oxygen (O)Atomic number = 8

Hydrogen (H)Atomic number = 1

Electron Shell Significance

• Electrons determine how an atom behaves when it encounters other atoms

• Outer orbital (valence shell) determines reactivity of atom - Electronegativity

• Atoms “desire” full outer orbitals

– Give up electrons (Na)

– Take electrons (Cl)

– Share electrons (O2)

• Noble gases - full outer shells (inert)

• Chemical reactions enable atoms to give up or acquire electrons in order to complete their outer shells

– These interactions usually result in atoms staying close together

– The atoms are held together by chemical bonds

Chemical Bonding and Molecules

Kinetic Theory Of Matter• All atoms and molecules are in constant random motion.

(Energy of motion is called kinetic energy.)

• The higher the temperature, the faster the atoms and molecules move.

• All motion theoretically stops at absolute zero.

Energy

• Energy is the capacity to do work or ability to cause change. Any change in the universe requires energy. Energy comes in 2 forms:

– Potential energy is stored energy. No change is currently taking place

– Kinetic energy is currently causing change. This always involves some type of motion.

Forms Of Energy

• Kinetic energy is the energy associated with motion

• Potential energy

– Is stored in the location of matter

– Includes chemical energy stored in molecular structure

• Energy can be converted from one form to another

• First Law Of Thermodynamics states that energy cannot be created or destroyed; energy can be transferred or transformed

On the platform, a diverhas more potential energy.

Diving converts potentialenergy to kinetic energy.

Climbing up converts kinetic

energy of muscle movement

to potential energy.

In the water, a diver has less potential energy.

Temperature, Pressure, And Volume

• Volume – Pressure Relationship

– At a constant temperature, volume is inversely proportional to pressure

• Volume – Temperature Relationship

– At constant pressure, the volume is directly proportional to temperature

• Cells constantly rearrange molecules by breaking existing chemical bonds and forming new ones

• Such changes in the chemical composition of matter are called chemical reactions

• Chemical reactions enable atoms to give up or acquire electrons in order to complete their outer shells

– These interactions usually result in atoms staying close together

– The atoms are held together by chemical bonds

• Reactions can be written as equations

Chemical Reactions

Reactants

Chemical Equations

• The chemical equation for the formation of water can be visualized as two hydrogen molecules reacting with one oxygen molecule to form two water molecules:

• 2H2 + O2 2H2O

Hydrogen gas Oxygen gas Water

Reactants Products

Reading Chemical Equations• The plus sign (+) means “react” and the arrow points towards the

substance produce in the reaction

• The chemical formulas on the right side of the equation are called reactants and after the arrow are called product

• The numbers in front of the molecules or atoms indicate the number of individual molecules or atoms (stoichiometric coefficients)

• The numbers behind are subscripts indicating the molecules or atoms are bonded

2Na + 2H2O 2NaOH + H2

Reactants Products

Chemical Reactions

• Are dependent on :

– Concentration

– Speed

– Energy (energy of activation)

– Orientation

Types Of Chemical Reactions

• Synthesis reactions - atoms or molecules combine to form a product

• Decomposition reactions - molecules breakdown into smaller molecules or atoms

• Exchange reactions - molecules exchange constituent components (swap partners)

• Reversible reactions - the product of a previous reaction can revert to the original reactants.

Combination (Synthesis) Reactions

• Combination (Synthesis) reaction occurs when two or more substances react to form products:

Na + Cl NaCl

Ca + 2NaCl CaCl2

• In both cases, Sodium and Calcium combine with Chlorine

Decomposition Reactions

• Decomposition reaction is when one substance undergoes a reaction to produce two or more substances:

2H2O 2H2 + O2

H2O2 H2O + O2

Exchange Reactions

• Exchange reaction occurs when molecules “swap partners”:

NaOH + HCl NaCl + H2O

H2CO3 + NaOH H2O + NaHCO3

Reversible Reactions

• Reversible reactions can go forwards (decomposition) or backwards (combination):

H2CO3 CO2 + H2O

• Chemical Equilibrium is defined as the state of dynamic balance in which the rates of forward and reverse processes (reactions) are equal

Chemical Products• Element: a substance composed of only one type of atom

(all the atoms have the same number of protons).

• Molecule: a unit composed of two or more atoms joined together by chemical bonds

• Compound: a substance composed of 2 or more elements that have been joined by chemical bonds

• Mixture: a combination of 2 or more substances that do NOT chemically bond e.g. sugar mixed with salt

Ionic Bonds• Atoms sometimes strip electrons from their bonding partners

• An example is the transfer of an electron from sodium to chlorine

• After the transfer of an electron, both atoms have charges

• A charged atom (or molecule) is called an ion

– An anion is a negatively charged ion

– A cation is a positively charged ion

• An ionic bond is an attraction between an anion and a cation - oppositely charged ions

• When an atom or molecule loses electrons, it becomes positively charged.– For example, when Na loses an electron it becomes Na+.

• Positively charged ions are called cations.• When an atom or molecule gains electrons, it becomes

negatively charged.• For example when Cl gains an electron it becomes Cl.

• Negatively charged ions are called anions.• An atom or molecule can lose more than one electron.• When molecules loose electrons, polyatomic ions are

formed.

Ions And Ionic Compounds

Ionic Compounds

• Compounds formed by ionic bonds are called ionic compounds, or salts

• Salts, such as sodium chloride (table salt), are often found in nature as crystals

Na+

Cl–

Covalent Bonds• Molecules are formed by covalent bonds

• A covalent bond is when two atoms share one or more pairs of outer-shell electrons (valence electrons)

• In a covalent bond, the shared electrons count as part of each atom’s valence shell

• Much stronger than ionic bonds – holds lots of Energy

• A single covalent bond, or single bond, is the sharing of one pair of valence electrons

• A double covalent bond, or double bond, is the sharing of two pairs of valence electrons

• Covalent bonds can form between atoms of the same element or atoms of different elements

Hydrogen atoms (2 H)

Hydrogenmolecule (H2)

Covalent Bonds

Oxygen (O2)

Name(molecularformula)

Electron-shell

diagram

Structuralformula

Space-fillingmodel

Covalent Bonds

Methane (CH4)

Name(molecularformula)

Electron-shell

diagram

Structuralformula

Space-fillingmodel

Water (H2O)

Name(molecularformula)

Electron-shell

diagram

Structuralformula

Space-fillingmodel

Covalent Bonds

Figure 2.9

• Outer orbital (valence shell) determines reactivity of atom - Electronegativity

• Electronegativity is an atom’s attraction for the electrons in a covalent bond

• The more electronegative an atom, the more strongly it pulls shared electrons toward itself

Electronegativity

H

O

H

H2O+ +

–

Polar Covalent Bond

• In a nonpolar covalent bond, the atoms share the electron equally

• In a polar covalent bond, one atom is more electronegative, and the atoms do not share the electron equally

The Structure Of Water– Its two hydrogen atoms are joined to one oxygen atom by

single covalent bonds

– But the electrons of the covalent bonds are not shared equally between oxygen and hydrogen

– This unequal sharing makes water a polar molecule

Unnumbered Figure 2.2

() ()

() ()

Hydrogen Bonds

• A hydrogen bond forms when a hydrogen atom covalently bonded to one electronegative atom is also attracted to another electronegative atom

• In living cells, the electronegative partners are usually oxygen or nitrogen atoms (b)

()Hydrogen bond

()

()()

()

()

()()

Figure 2.11b

–

Water(H2O)

Ammonia(NH3)

Hydrogen bond

+

+

–

+

+

+

Hydrogen Bonds

Weak Chemical Bonds

• Most of the strongest bonds in organisms are covalent bonds that form a cell’s molecules

• Weak chemical bonds, such as ionic bonds and hydrogen bonds, are also important

• Weak chemical bonds reinforce shapes of large molecules and help molecules adhere to each other

Biological Importance Of Water• Acts as a powerful solvent

• Participates in chemical reactions

• Water has a high specific heat which moderates temperature - absorbs and releases heat very slowly, minimizes temperature fluctuations to within limits that permit life

– Heat is absorbed when hydrogen bonds break

– Heat is released when hydrogen bonds form

• Requires a great amount of heat to change to a gas

– Heat of vaporization - the quantity of heat a liquid must absorb for 1 gram of it to be converted from a liquid to a gas

– Evaporative cooling - Allows water to cool a surface due to water’s high heat of vaporization

• Acts as a lubricant

Polarity & Hydrogen Bonds• Cohesion - molecules attract other

water molecules

• Capillarity

– Water molecules are drawn up a narrow tube

– Helps pull water up through the microscopic vessels of plants

• Surface tension

– water molecules on the surface cling to each other – related to cohesion

– Is a measure of how hard it is to break the surface of a liquid

• Adhesion - water molecules attract other charged substances

Water As A Solvent• Water is a versatile solvent due to its polarity

• It can form aqueous solutions

• The different regions of the polar water molecule can interact with ionic compounds called solutes and dissolve them

Negative oxygen regions

of polar water molecules are attracted to sodium

cations (Na+).+

+

+

+Cl –

–

–

–

–

Na+

Positive hydrogen regions

of water molecules cling to chloride anions

(Cl–).

++

+

+

–

–

–

–

–

–Na+

Cl–

Formula and Molecular Weights

• Formula weights (FW) is the sum of the atomic weights of each atom in the chemical formula.

• FW (H2SO4) = 2AW(H) + AW(S) + 4AW(O)

• = 2(1.0 amu) + (32.0 amu) + 4(16.0)

• = 98.0 amu

• If the chemical formula is also its molecular formula then the weight is called the molecular weight (MW).

• MW(C6H12O6) = 6(12.0 amu) + 12(1.0 amu) + 6(16.0 amu) =????

The Mole• The unit we use to express the quantity of atoms, ions, and

molecules that an object contains is called mole.

– Mole: convenient measure chemical quantities.

• The actual number of atoms, ions, or molecules in 1 mole of something = 6.0221367 1023 (Advogadro’s number).

• Thus,

• 1 mole of 12C atoms = 6.02 x 1023 12C atoms

• 1 mole of H2O molecules = 6.02 x 1023 molecules

• 1 mole of NO3- ions = 6.02 x 1023 ions

Visualizing The Mole Concept

Different Units

Solution Composition• Solutions are homogenous mixtures of two or more substances:

– Solute: present in smallest amount and is the substance dissolved in the solvent.

– Solvent: present in the greater quantities and is used to dissolve the solute.

– Example: NaCl dissolved in Water (water = Solvent and NaCl = Solute)

• Change concentration by using different amounts of solute and solvent.

• Molarity: Moles of solute per liter of solution.

Concentration of Solutions • Percent solutions

• Ratio of solute to solvent expressed as a percentage: weight (g)/volume (ml)

• Unit seen on IV bags and medicinal solutions

– 5% dextrose = 5g dextrose / 100 ml of solution

– 0.9% saline = 0.9g NaCl / 100 ml of solution

Example

• Betadine antiseptic solution contains 10g of povidine-iodine in 100mL of solution. Calculate the percent (w/v) concentration of the solution.

% w/v = grams of solute x 100mL of solution

= 10g/100mL x 100= 10%

Formula for Molarity

• The most widely used way of quantifying concentration of solutions in chemistry. Molarity is generally represented by the symbol M and defined as the number of moles of solute dissolved in a liter of solution.

Concentrations Of Solutions

General Properties Of Aqueous Solutions (Terms)

• Acids - substances that able to ionize in solution to form hydrogen ion (H+) and increase the concentration of H+ in the solution.

• For example, HCl dissociate in water to form H+ and Cl- ions.

• Bases - are substances that can react with or accept H+ ions.

• For example, OH- will accept H+ from HCl forming H2O.

• Salts - are ionic compounds that can be formed by replacing one or more of the hydrogen ions of an acid by a different positive ion.

• For example, NaCl instead of HCl.

Acids, Bases, And pH

• Dissociation of water molecules leads to acidic and basic conditions that affect living organisms

– Water dissociates into hydronium ions and hydroxide ions

– Changes in the concentration of these ions can have a great affect on living organisms

H

Hydroniumion (H3O+)

H

Hydroxideion (OH–)

H

H

H

H

H

H

+ –

+

Acids, Bases, And pH• Acid - A chemical compound that dissociates into

one or more hydrogen ions (H+) and one or more negative ions (anions). An acid donates H+ ions (protons) to solutions

• Base - Dissociates into one or more positive ions (cations) and one or more hydroxide ions (OH-). A base accepts H+ ions and removes them from solution, reducing the hydrogen ion concentration of a solution

Strong And Weak Acids

• Strong Acids ionize or dissociate completely in water

HCl + H2O H+(aq) + Cl-

(aq)

– Strong Acids: HNO3 (nitric acid); H2SO4 (sulfuric acid)

• Weak Acids ionize slightly (less than 5%) in water

– CH3COOH (acetic acid)

– H2CO3 (carbonic acid)

– H3PO4 (phosphoric acid)

• To describe the acidity of a solution, we use the pH scale

• Is a measure of the concentration of H+ ions in a solution

• Is determined by the relative concentrations of H+

– Basic = High pH = few H+, many OH-

– Acidic = Low pH = many H+, few OH-

Basicsolution

Neutralsolution

Acidicsolution

Oven cleaner

Household bleach

Household ammonia

Milk of magnesia

Seawater

Human bloodPure water

Urine

Tomato juice

Grapefruit juice

Lemon juice;gastric juice

pH scale Figure 2.17

pH

pH Scale• The pH scale is a logarithmic scale

used to express the amount of H+ ions in a solution.

• pH is defined as the negative log of the [H+] of a solution

• A change of one whole number represents a tenfold (10X) change in the number of H+ ions.

• A solution with pH 3 has ten times as many H+ ions as a solution with pH 4

Buffers• A buffer is a substance that helps minimize the change in the pH of a

solution when acids or bases are added.

• Consist of an acid-base pair that reversibly combines with hydrogen ions

• Buffers work by releasing H+ when their concentration falls, and absorbing H+ when their concentration rises.

• Buffers are important to living organisms because most cells can survive and function normally only within a relatively narrow range

-

pH

[H+]

1

14

Bicarbonate Buffer System• A mixture of carbonic acid (H2CO3) and its salt, sodium bicarbonate (NaHCO3)

• If strong acid is added:

– Hydrogen ions released combine with the bicarbonate ions and form carbonic acid (a weak acid)

– The pH of the solution decreases only slightly

• If strong base is added:

– It reacts with the carbonic acid to form sodium bicarbonate (a weak base)

– The pH of the solution rises only slightly