Chapter 2 chemistry part 1 inorganic compounds

matter

has mass and takes up space

atom

smallest stable unit of matter

molecules

when atoms combine (two or more held by chemical means)

element

all of the same type of atom, cannot be broken down into simpler  substances by ordinary means. (note:  periodic table of elements, 118 known)

compound

2 or more elements held by chemical bonds

Example:  (glucose molecule) C₆H₁₂O₆ is a compound,   O₂ is not

II. Energy (E)

the capacity to do work

potential

the E an object has because of its position in relation to other objects.

  (i.e. an object has the capability to do work) =   “stored energy”

kinetic

the E associated with a moving object (the form of E that is actually doing work) = “energy of motion”

III. Atomic structure

Fundamental particles: protons

(+) in nucleus and one mass unit *protons and neutrons similar in size and mass

neutrons

(no charge) in nucleus and one mass unit *protons and neutrons similar in size and mass

electrons

(-), smaller (1/1800^th the size),

considered weightless

located in orbitals/clouds surrounding the nucleus

IV atomic measurements

periodic table

lists measurements, appendix E (94 naturally occurring)

1.atomic number

the number of protons in nucleus is written to the left of the symbol   *(Note:  since the number of protons in a NEUTRAL atom always equals the number of electrons, the atomic number indirectly tells the number of electrons.)

2.mass number

sum of the masses of its protons and neutrons (mass of electrons is so small that it is ignored)

atomic weight

derived from an average mass, determined for any individual element.

V isotopes

radioisotopes

elements with unstable nuclei that emit subatomic particles in measurable amounts. (alpha, beta, or gamma particles)

Half life

time required for 50% reduction in radioactivity

(seconds/hours to tens of thousands of years)

-Used in Medical Imaging; measures vitamin absorption, blood volume, organ functioning…etc. (uses isotopes with short half-lives)

VI. 4 Elements of the human body:

VII. Matter in Combination

A, mixture

:  2 or more substances physically intermixed (not bonded)

Three types

1.solution

when two or more substance are   EQUALLY mixed

  “homogeneous”

solvent

dissolving compound (largest quantity)

solute

compound being dissolved (smallest   quant.)

colloid – (ex. jello, cytosol in cells)

when two or more substances are UNEQUALLY  

  mixed, “heterogeneous”

  - particles do not settle out, solutes generally larger.

3.suspension (ex. blood)

Also, a heterogeneous mixture

  - Large, sometimes visible solutes

  - Solutes will settle out if left long enough

  (Ex. antibiotic suspension, sand and water)

Principles of chemical bonds

I. Atomic Energy Levels

Electrons occupy specific areas in the atom, the amount of E that can be contained in any one space (level) determines the number of electrons that will be present at the specific level.

Energy Levels (shells):

Rule of 8 will govern reactivity

“octet rule”

Ex. Ne, an inert (Noble) gas, has a stable configuration of   electrons in its “outer shell”, also called its   “valence” shell.

Ex. Na, a highly reactive element, will tend to lose the outer electron to get down to the most stable configuration (8 in the outer shell= “valence shell”)

II. atomic interactions

A. Two Rules Govern Interactions:

1. An unfilled electron shell is “unstable”

  Atoms tend to interact to “fill” the electron shell via   gaining, sharing, or   losing electrons. (full shells: #1 = 2   electrons, #2 & 3 = 8 electrons)

  2. Number of electrons in outer shell   determines properties of an element

  (Periodic Table helps predict reactivity)

Ions defined

atoms or molecules that have a + or - charge

cation

positive chg. ion (loss of electrons)

anion

negative chg. ion (gain of electrons)

Ionic bonds

•Atoms in this type of bond donate or take on electrons

•Results in a stable outer shell

•Occurs between particles that are charged (ions)

b. Types of molecular bonds:

1.Ionic bonds

2.non polar covalent bonds

Ex. H₂, Hydrogen molecule, single covalent bond

  Goal:  to attain a stable electron configuration   (i.e. to fill the energy level requirements)

2.Non-polar Covalent Bonds

Goal:  to attain a stable electron configuration   (i.e. to fill the energy level requirements)

  Ex. O₂, Oxygen molecule, double bond

  Ex. CO₂, Carbon dioxide, double bond

  Ex. N₂, Nitrogen molecule, triple bond

Characteristics of covalent bonds

strong

often electrically neutral (equal sharing of electrons)

If non-polar indicates no “poles” to molecule

3. Polar covalent bonds

Sharing of electrons is UNEQUAL

Ex. water H₂O

Neg. charge predominates at oxygen

“Separation of charges exists”

4. Hydrogen Bonds (attraction)

After forming polar or ionic bonds with other elements (oxygen and nitrogen frequently)…

The weakly positive hydrogen atoms can be attracted to nearby atoms/ions that carry a negative charge.

Important for inter-molecular attractions - Ex. 3-D

surface tension in water, or structures of proteins

III. chemical notation

Used to describe events in a precise and concise way.

A. Rules

1. Element symbol indicates one atom (know these):

Ex. H, C, N, O, Fe, Mg, Zn, Ca, Na, K, P, CI, S, I

2. The number preceding the symbol indicates more than one atom/molecules.

Ex. 2 He, there are two helium atoms

Ex. 4 H₂0, there are four water molecules

3. Number (subscript) written following symbol indicates the number of that specific element in a molecule.

Ex. H₂ = one hydrogen molecule (2 H atoms bound   together

  Ex. 3 H₂ = 3 hydrogen molecules (6 H atoms)

  Note:  NO subscript means one atom present in   molecule

  Ex. Water, H₂O

4. writing reactions (rxn), (shorthand)

5. All RXNs must be balanced-

Reactants must equal the Products

  (No. of atoms on one side of the arrow equal the no. of   atoms on the other side of the arrow)

2 H    +    O  —)  H₂O ,   balanced but Hydrogen and Oxygen atoms don’t exist in nature by themselves, so must change the equation.

2 H₂    +    O₂ —)  H₂O ,   formulae are accurate, but   equation is not balanced

2 H₂    +    O₂ —) 2 H₂O  ,  Balanced with correct   formulae

IV. chemical reaction

A. characteristics of chemical bonds

1. contain energy

2. stronger bond means more E: (see textbook)

 STRONGEST ^ covalent bonds (nonpolar)

Covalent bonds (polar)

Ionic bonds

WEAKEST v Hydrogen bonds

IV. chemical categories

B. categories of reactions

1. decomposition -

molecules broken down into   smaller fragments

AB   —) A  +  B   +  (E)   * Used by organisms to do work.

E = energy available

2. synthesis

smaller fragments are assembled to   make larger   molecules.

 

  (E)  +   A   +   B   —)  AB               E = energy used

IV. chemical reactions

B. categories of reactions

3. exchange -

reactants and products contain the same   components in different combinations (aka displacement reactions)

  AB   +   CD   —)  AC    +    BD

a.  Exergonic rxn –

when E released from the   decomposition rxn is more than the E required for   synthesis rxn.

b.  Endergonic rxn

when E required for synthesis is more than E produced from decomposition rxn.

IV. chemical reactions

C. characteristics of chemical reactions

1. reversible (theoretically)

IV. Chemical Reactions

     C.  Characteristics of Chemical Reactions

  2. Chemical Equilibrium–

at some point the product AB is synthesized at the same rate AB is

decomposed to A  +  B.

  “AT EQUILIBRIUM”

     3.  Factors may influence rate of a reaction