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2.1: Atomic Theory of Matter

400 B.C. -Democritus: matter made of tiny, indivisible spheres called atoms

1800s- John Dalton: Atomic Theory

1.elements made of indivisible atoms

2.all atoms of 1 element have identical properties unique to the element

3.atoms can’t be created, destroyed or changed

4.compounds result from the combination of atoms in small, whole number ratios

5.relative #s and kinds of atoms in a compound are constant

ATOM: the smallest part of an element that maintains its chemical identity chemical and physical changes

      Made of:

            1. Protons: in nucleus, + charge (at. #)

            2. Neutrons: in nucleus, 0 charge

            3. Electrons: around nucleus, - charge

      Neutral atoms have the same # of p⁺ and e^-

MOLECULE: smallest part of an element/compound that can have a stable, independent existence

      Diatomic molecules: O₂, N₂, H₂, F₂, Cl₂, Br₂, I₂

      Polyatomic molecules: P₄, S₈

CHEMICAL FORMULA:

~shows chemical composition of a substance

~represents elements present in their proper ratios

LAW OF CONSTANT COMPOSITION (definite proportions):

Different pure samples of a compound always contain the same elements in the same proportion by mass; atoms of these elements combine in fixed numerical ratios EX: NaCl, MgCl₂, NaOH

LAW OF MULTIPLE PROPORTIONS: When 2 elements A and B form more than 1 compound, the ratio of the masses of B that combine with a given mass of A can be expressed by small whole numbers.

i.e. elements combine in defined, small, whole number ratios

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Example problem:

Phosphorus forms 2 chlorides.

compound A: 4.35g P combines with 25.65g Cl

compound B: 6.61g P combines with 23.39g Cl

we can develop a whole # ratio of the masses of Cl to P for each of the compounds, then simplify it to small, whole numbers. So...

A: 25.65g Cl/4.35g P = 5.90 g Cl/g P

B: 23.39g Cl/6.61g P = 3.54 g Cl/g P

examining these ratios more carefully,

5.90: 3.54 simplifies to 1.67: 1

the ratio of the masses needs to be in whole #s, so this can be multiplied by 3 to equal 5: 3.

This 5:3 ratio is the small, whole # ratio of masses that is discussed in the law of multiple proportions.

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2.2: The Discovery of Atomic Structure

The most convincing evidence for the existence of electrons came from cathode ray tube experiments

-They are glass tubes which contain gas at a low pressure

-Two electrodes (two pieces of metal) at the end have been sealed

cathode - negative electrode

anode - positive electrode

-When high voltage is applied, current flows from the cathode to the anode

-The electron beam usually travels in a straight line but is deflected if a magnetic field or an electric field are present

-Because the beam is produced at the negative electrode toward a positive plate, it was proposed that is made up of very tiny, negatively charged particles

In 1897, Thompson studied the negatively charged particles more carefully and called them electrons

J.J. Thomson determined:

1. The ratio of charge to mass for electrons

2. That all atoms had a fundamental particle called an electron

The "Plum Pudding" Model

Further experiments were conducted to determine either the charge or the mass of an electron

Millikan’s Oil Drop Experiment

Experiment:

-A fine mist of oil was sprayed into a box and the oil droplets were allowed to fall between two horizontal plates

-Falling oil droplets has a negative charge so they were suspended between two electrically charged plates

-The mass was determined by how rapidly the oil drops fell through the air

-Once the mass was known, the charge was determined

Millikan Determined:

1. The charge of an electron

2. That atoms contain a number of electrons

Canal Rays and Protons (FYI)

Eugene Goldstein observed in 1886 that a cathode ray tube also generates canal rays

A canal ray is a stream of positively charged particles (cations) that moves toward the cathode in a cathode ray tube

The positive rays are created when cathode rays knock electrons from the gaseous atoms in the tube, forming positive ions

atom -------------> cation⁺ + e^-

Goldstein also found that different elements give positive ions with different charge to mass ratios

The regularity of the charge to mass ratio led to the idea of a unit of positive charge which resides in the proton

A proton is a fundamental particle of the atom with a charge equal in magnitude but opposite in sign to the electron

A proton is 1836 times the mass of an electron

RUTHERFORD AND THE NUCLEAR ATOM

Refer to diagram 5-4 on Pg. 160

Rutherford shot alpha (positive) particles at a this sheet of gold foil

he expected all of the particles to go straight through the foil but...

    * Most went straight through

    * There were a few slight changes in the angle at which the particles hit the screen behind the foil

    * There were 1 in 8000 deflected right back at the alpha source

 "It was almost as if you fired a 18-inch shell into a piece of tissue paper and it came back and hit you"

What does this mean???

* There is a positive portion to the atom (like charges repel)

* It takes up a tiny fraction of the volume of the atom (very few wide angle deflections)

* It makes up most of the mass of the atom (the positive thing is small and dense)

* The atom has a nucleus!!!

2.3: The Modern View of Atomic Structure

Atomic Number

A few years after Rutherford's scattering experiments, H.G.J. Moseley studied X-rays that are given off by certain elements

He found that the wavelengths of the X-rays emitted by an element are related in a precise way to the atomic number of the element

On the basis of his mathematical analysis of these X-ray data, he concluded that:

~ Each element differs from the preceding element by having one more positive charge in its nucleus

~ The number of protons in the nucleus of an atom determines its identity

~ The atomic number of an element is the number of protons in the nucleus of an atom

Mass Number and Isotopes

Isotopes are atoms of the same element with different masses. They have the same number of protons but different numbers of neutrons

Example of the different isotopes of hydrogen:

1. Hydrogen (hydrogen-1) = no neutrons
2. Deuterium (hydrogen-2) = 1 neutron
3. Tritium (hydrogen-3) = 2 neutrons

Each of these contain one proton in the atomic nucleus

*All three isotopes of hydrogen display very similar chemical properties

Mass number = number of protons + number of neutrons

= atomic number + neutron number

The mass number for normal hydrogen atoms is 1, for deuterium is 2 and for tritium is 3

Nuclide symbol = the symbol for an atom

                            E = symbol for the element

                            Z = atomic number

                            A = mass number

The Atomic Weight Scale and Atomic Weights

The atomic weight scale is based on the mass of the carbon-12 isotope

*One amu (atomic mass unit) is exactly 1/12 of the mass of a carbon-12 atom

For example, the mass of one atom of C is exactly 12 amu

Important:

1 gram = 6.022 x 10²³ amu or 1 amu = 1.660 x 10^-24g

Important Terms:

1. atomic number: (Z) is equal to the number of protons in the nucleus and the number of electrons in a neutral atom

2. mass number: (A) is equal to the sum of the number of protons and the number of neutrons in the nucleus of an atom of a particular isotope

3. atomic weight: is the weighted average of the masses of the isotopes

Calculating Atomic Weight

Example:

 Three isotopes of magnesium occur in nature. Their abundances and masses were determined by mass spectrometry and are given below:

Use this information to calculate the atomic weight of magnesium.

 2.5: The Periodic Table

Know the following:

2.6: Molecules and Molecular Compounds

MOLECULE: smallest part of an element/compound that can have a stable, independent existence

      Diatomic molecules: O₂, N₂, H₂, F₂, Cl₂, Br₂, I₂

      Polyatomic molecules: P₄, S₈

MOLECULAR COMPOUNDS: formed from molecules that can exist individually

      IONIC COMPOUNDS: formed from + and – charged ions

                                ~ions can’t exist separately

                                ~ these exist as formula units (FU)

EMPIRICAL FORMULA: the smallest whole # ratio of atoms present in a compound

MOLECULAR FORMULA: the true formula of a compound that is either the same as or an integer multiple of the simplest formula

2.7: Ions and Ionic Compounds

ALL ABOUT IONS:

~ion: atom/group of atoms with an electrical charge

~cation: positively charged ion (metals lose e^-)

~anion: negatively charged ion (nonmetals gain e^-)

~polyatomic ion: group of atoms with an electrical charge

know the ion charges on the common ion chart!!

WHY ARE THERE IONS?

All atoms are wannabes: they wanna be stable (gr.18)

Form ions

EX: Na= 1s² 2s² 2p⁶ 3s¹          Na⁺= 1s² 2s² 2p⁶

Cl = 1s² 2s² 2p⁶ 3s² 3p⁵           Cl^- = 1s² 2s² 2p⁶ 3s² 3p⁶

Ions achieve noble gas configuration (8e-)

Refer to periodic table

FORMING A COMPOUND: NaCl

      Na+ and Cl-

      +1 + -1 = 0 (no net charge)

WRITING FORMULAS:

1.cation first

2.anion second

3.ions in a ratio with no net charge (cross & drop)

extras: roman numerals, simplify subscripts, polys

PRACTICE: sodium fluoride, silver bromide, calcium oxide, zinc sulfide, iron (II) bromide, copper (II) oxide, iron (III) sulfate, copper (I) phosphate, aluminum phosphate

NAMING COMPOUNDS:

1. ID the ions

2. Name of cation first

   3. Name the anion

extras: roman numerals

PRACTICE: NH₄Br, ZnF₂, Al(CH₃COO)₃, Cu₂CO₃, Fe(OH)₃, FePO₄, MgSO₃, Ag₂O

2.8: Naming Inorganic Compounds

Molecular compounds are made of 2 (or more) nonmetals

Naming:

Use prefixes to indicate the # of each atom (except for single atoms)

Writing formulas:

~Use the prefix in the name as the subscript for that element

~Don't simplify the subscripts

Special Cases: Acids

1. H + single anion: binary acid named as "Hydro___ic acid"

    ~ ___ is the root of the nonmetal

2. H + polyatomic ion: oxyacid

    a. If the anion ends in "-ate", the acid name is "___ic acid"

    b. If the anion ends in "-ite", the acid name is "___ous acid"

        ~ ___ is the root of the nonmetal (see chart in text)

2.9: Simple Organic Compounds

Hydrocarbons: Compounds made of carbon and hydrogen

ALKANES: end in -ane

If H is replaced by -OH, the molecule is an ALCOHOL (also named according to the # of C's); ends in -ol

        Replacing an H on the 1st C is indicated in the name: 1-propanol

        Replacing an H on the 2nd C is also indicated in the name: 2-propanol

The presence of a double bond indicates an ALKENE

a C=O bond indicates a carboxylic acid, ketone, aldehyde or ester