The Nature of Matter

Chapter 1BIOLOGY 391

What is everything made of?

• MATTER Anything that has mass and takes up space

• ATOM The smallest unit of matter

ATOMS• Basic unit of matter

• Size: 1,000,000 (million) side by side = 1 cm

• Atoms like to be neutral- no charge

– Equal number of protons and electrons

– There are specific numbers of “sub-atomic” particles that the atom wants• Special cases are called isotopes or ions

STRUCTURE OF ATOM

– NUCLEUS - protons and neutrons held together by the “strong force”

• Protons (+)• Neutrons (o)

– ELECTRON CLOUD (orbitals) – electrons surround nucleus • Electrons (-) only contains about 1/1836 the

mass of proton or neutron• Constantly moving within orbital- attracted to

the nucleus by the “weak force”

ATOMS •The smallest particle of an element that can exist and still have the properties of that element.•An atom has many subatomic particles

Particle Charge Location Mass

Proton + Nucleus 1 amu *

Neutron N Nucleus 1 amu

Electron - ElectronCloud

0 **

•Atomic mass unit or Dalton

•** The mass of an e- is so small it’s negligible.

Electron orbitals

1st orbital can only hold 2 electrons (too close to nucleus- not much space)

2nd orbital can hold up to 8

3rd orbital can hold up to 8

How many atoms are there?PERIODIC TABLE

Atoms are grouped as Elements

• Are listed on the Periodic Table• Dimitri Mendele’ev 1869• Arranged in order of their atomic #’s• Table is divided into Groups and Periods• The atomic number and atomic mass are

given for each element. Example: 6

C12.01

Atomic #

Atomic Mass

Element

“Groups”Vertical Column

Numbered 1-8All elements of same group have the same # of e-’s in their valence shell & have similar chemical properties

Atomic # increases by 1 from left to right Example:

H, Li, Na, K, Rb, Cs, FrAll have 1 e- in valence (outer) shell

“Periods”• Is a horizontal row of elements• There are 7 periods• Elements in a period have the same #

of energy levels

Example:H and He are in period 1They have 1 energy level

The Elements are often described as Metallic and Nonmetallic

Atomic Number# of protons

(and also # of electrons)

Chemical symbol

Name of Element

Atomic MassThe weight Of carbon

atom oraverage

weight of all isotopes

6

CCarbon

12.011

What is the difference between atoms?

Special Form: ISOTOPE

Atoms of the same element containing different numbers of neutrons in the nucleus

– Some give off radiation – used to:• Trace atoms through a reaction or an organism• Treat cancer• Date very old, once living organisms

2 examples of Isotopes

Recap• What are the three subatomic particles and

their charges?• What is the only actual difference between

gold and mercury? • What is the atomic mass of lead?• What is an isotope?

Organization of Matter

• Atoms usually do not occur alone, but exist with other atoms as:

–Elements (all of the same atoms)–Molecules or Compounds

• Same or different atoms bonded

MATTER: anything that has mass & takes up space

ProtonsNeutronsElectrons

Does not contain C-C bonds

LIFE!

Organic

Molecules orCompounds

Inorganic

C-C, C-H bond

ATOMSElements: Shown in

Periodic Tablepure

bonded

Elements

• A substance which cannot be split into simpler substances by a chemical rxn.

• A grouping of the same type of atoms– ORDER MATTERS!

• More than 100 elements exist (shown in the periodic table)

• Carbon Elements:

Elements of Life

92 naturally occurring elements

Elements Found in Living OrganismsN CHOPS (macronutrients)

C HOPKINS Ca Fe Mg B Mn Cu Cl Mo Zn

MATTER: has mass & takes up space

ProtonsNeutronsElectrons

Does not contain C-C bonds

LIFE!

Organic

Molecules orCompounds

Inorganic

C-C, C-H bond

ATOMSElements: Shown in

Periodic Tablepure

bonded

Why do atoms bond?• An atom wants to have a complete outer shell

of electrons.

To do this, it can…• Share electrons with another atom• Give away its extra electrons• Steal extra electrons

• Bonds store ENERGY

Remember: an atom is when its outer orbital is filled

For the “smaller” elements, the outer shell holds 8. The electron clouds Increase in size as you go across the periodic table. HOWEVER, the trend stays the same for the number needed to complete the outermost shell.

BONDING: depends on number of electrons in the outermost orbital

shell• Covalent Bond- atoms share a pair of electrons

sometimes share 2 (double bond) or 3 (triple bond) pairs

Ionic Bond- One atom (very unstable) gives 1, 2 or 3 electrons away to another atom. The atom that loses electrons becomes positively charged. The atom that gains the electrons becomes negatively charged. The opposite charges cause the atoms to “bond” together

(opposites attract).

Molecules bonded atoms

“Molecule” is often used to refer to an individual grouping. “singular”

Compounds: also, 2 or more atoms bonded together.

Often referred to as a larger conglomerate of bonded molecules

• Order matters!• These 2 cmpnds are

made of the same 3 atoms, but in a different arrangement.

• The arrangement is responsible for their affects.aspirin sucrose

Bonding• IONIC• COVALENT

– Polar– Nonpolar

• VAN DER WAALS• HYDROGEN

The periodic table shows the pattern of electrons each element has in its

outermost shell

Ions • Atoms that have lost or gained electrons

cation – positively charged ion – lost electronsanion – negatively charged ion – gained electrons

• Ionic bonds are weaker than covalent bonds– Hold less energy in the bond

Cations versus Anions• Atom that loses electrons• Positively charged ion• Elements in Groups 1, 2,

and 3 tend to lose electrons• Metallic elements tend to

form positive ionsExample:

Ca Lose 2 electrons+2 charge

• Atom that gains electrons• Negatively charged ion• Elements in groups 5, 6, and

7 tend to gain electrons• Nonmetallic elements tend

to form anionsExample:

ClGain 1 electron-1 charge

IONIC BONDING

Na (sodium) is very unstable because it only has one e- in its outer orbital. Cl’s (chlorine) outer orbital is

almost filled. Na gives its lonely e- to Cl.

Na become Na+ Cl becomes Cl-

Their opposite charges cause them to be attracted to one another- This is an ionic bond.

Ionic Compounds• Metals react with nonmetals forming ionic compounds• Salts• Held together by electrostatic forces

Example: + attracted to –• Most are crystalline solids at room temperature• When dissolved in water they conduct electricity

Ex. Sodium Chloride or Table SaltNa+ + Cl-

• Dissociate(break apart) in water, producing free ions

Electrolyte• A solution that conducts electricity.• Term for salts, specifically ions.

• The term electrolyte means that this ion is electrically-charged and moves to either a negative (cathode) or positive (anode) electrode

• Ions that move to the cathode (cations) are positively charged

• Ions that move to the anode (anions) are negatively charged

Covalent Bonds• Bonds formed when atoms share electrons• Atoms with 4 or 5 e-’s tend to share • Each e- spends part of its time around one

nucleus and then around the other nucleus.

• Sharing e-’s completes the valence shell

Example of a Covalent BondExample of Covalent Bonding- Water

Covalent Bonds• Bonds between non-metals

• Poor conductors of electricity

• Do not dissociate easily in water

• Two Types:Polar: Unequal SharingNon-Polar: Equal Sharing

Van der Waals forces• Attraction between oppositely charged areas

of adjacent molecules• Remember, e- are constantly swarming

around the nucleus. At times, there may be a moment of asymmetry – All of the electrons might be on one side– This sets up a “dipole”

weaker than covalent bonds and ionic bonds

RECAP• What is an ion?• Why is it important that atoms bond?• What causes atoms to bond?• Explain the difference between an ionic bond

and a covalent bond• What are Van der Waals forces?

BONDS

IONIC

Transfers/takes electronsPositive/Negative chargesWeak bondEx: NaCl (salt)

COVALENT

Shares electrons NeutralStrong bondEx: H2O, CO2, NH3

WATER – Why so Great?

Water!• What is the chemical formula for water?• How much water covers the Earth?• How much of your body is water?• Is there water in food?• How long could you live without water?

• H2O• 75% • 60-70%• Yes!• 3 days

Properties of Water

• Phases: Solid, Liquid, Gas• Polarity• Hydrogen bonds

– Adhesion– Cohesion

• Making Mixtures– Solutions– Suspensions

• Making Acids and Bases

Water Density

• Ice is less dense than liquid water• When water freezes air is trapped within the

frozen ice making the cube larger and less dense

• Benefits:– Fish and plant life can survive in liquid layers of

water under ice

PHASE CHANGES: the closeness and speed of the compounds

Polarity

• Water is polar

• Although the compound is neutral overall there is a shift of charge within the compound

The much larger atom, Oxygen, pulls more on the shared e-

This end of the compound becomes slightlymore negative.

Hydrogen ends become slightly

positive

Hydrogen Bonding• Due to polarity, water compounds attract to

one another• Slightly negative oxygen attracts slightly

positive hydrogen from another compound• This attraction among water is COHESION.• Water is also attracted to other materials.

This is ADHESION.

COHESION

Water compounds attractTo glass molecules

And form a meniscus

Water compounds attract To one another-

causes water to “bead”

ADHESION

The greatest solvent on Earth!

• Water’s polarity allows it to break ionic bonds of other compounds…creating free ions.

Mixtures Two or more elements physically mixed together but

not chemically combined (not bonded)

1. SOLUTIONS- a solute is dissolved into a solvent – Distributes evenly– “Like dissolves Like”– Ex: Koolaid, salt water

2. SUSPENSIONS- added substance does not dissolve but breaks into small enough pieces that it remains suspended in the water and does not settle out.- Ex: blood

RECAP

• Why does ice float on a lake?• Explain the polarity of water – how are the

charges distributed?• What is the difference between adhesion and

cohesion?• Explain the difference between a solution and

a suspension

Water Dissociation

• Water can break apart on its own into 2 charged ions– Hydrogen ions and hydroxide ions

Acids and Bases

Water can react to form individual ions:

H2O H+ + OH-

• In pure water this occurs naturally but the amount of H+ is always = to the amount of OH- so water remains neutral

pH scale: “the power of Hydrogen”

• Some solutions made with water become acidic or basic. This is determined by the amount of H+ (hydrogen ions) in the solution

• pH = - log [H+]– Example: [H+] = 1 x 10 -5 pH = 5 (acid)

[H+] = 1 x 10 -9 pH = 9 (base)Remember how exponents work:

0.00001 is greater than 0.000000001

• b/c it’s logarithmic, each pH unit represents a tenfold difference in concentration of H+ ions

Acids

• pH range from 0 – 6.99 • Any compound that forms H+ ions in solution• H+ ions > OH- ions• The closer to 0 the more acidic the solution• Examples: stomach acid, lemon juice

Bases (Alkaline)

• pH ranges from 7.01 to 14• Any compound that forms 0H- ions in solution• OH- ions > H+ ions• The closer to 14 the more basic the solution• Examples: lye, bleach, oven cleaner

ACID:Any compound that forms H+ ions in solution

BASE: Any cmpnd that forms 0H- ions in solution

pH and Living Things• pH values in living cells are usually kept

between 6.5 and 7.5– Optimal pH for chemical reactions to take place in

the body– Any switch in pH could cause serious/fatal

problems

Buffers

• Weak acids or bases that can react with strong acids or bases

• Used to regulate pH and prevent sharp sudden changes in pH

• There are natural buffers in your blood that keep the pH at 6.5 to7.5

RECAP

• What makes a solution acidic or basic?• How is acidity measured?• A solution with pH 8.5 is considered….

Forming…Acid Rain…Reacting with stone