PSMA (General Chemistry pt. 1)

Atoms

it consists of a positively charged core (the atomic nucleus) which contains protons and neutrons, and which maintains a number of electrons to balance the positive charge in the nucleus

Democritus (460-370 BC)

• Atomos

• aka Discontinuous matter

Plato and Aristotle

“there can be no ultimately indivisible particles”

Dalton’s Atomic Theory

• Elements are composed of extremely small particles, called atoms.

• All atoms of a given element are identical, having the same size, mass, and chemical properties. The atoms of one element are different from the atoms of all other elements.

• Compounds are composed of atoms of more than one element. In any compound, the ratio of the numbers of atoms of any two of the elements present is either an integer or a simple fraction.

• A chemical reaction involves only the separation, combination, or rearrangement of atoms; it does not result in their creation or destruction

Modern Atomic Theory

• Atoms may be disintegrated. In nuclear reactions, atoms are being transferred into atoms of single elements in a process known as nuclear transmutation.

• Not all atoms of any given element are alike

• Not all atoms of a given element pose identical properties except in mass

• Atoms of different elements have different properties

• Electrons

• Protons

• Neutrons

What are the structure of Atoms

Electrons

• Discovered by Joseph John Thomson

• was the first component of the atom to be identified

• a mass of 9.109 × 10-31 kg

• a charge of -1.602 × 10-19 coulombs

Protons

• One of the components of a nucleus

• discovered by Eugene Goldstein

• a mass of 1.673 × 10-27 kg

• a charge of +1.602 × 10-19 coulombs

Neutrons

• Another component of a nucleus

• discovered by James Chadwick in 1932

• a mass of 1.675 × 10-27 kg

• has no charge

John Dalton (1803)

History:

Solid sphere model

J.J. Thomson (1904)

History:

Plum pudding model/ "Raisin Bread Model”

Ernest Rutherford (1911)

History:

Nuclear model

Niels Bohr (1913)

History:

Planetary model

Erwin Schrodinger (1926)

History:

Quantum model

1. Thomson model

2. Rutherford model

3. Bohr model of an atom

4. Rutherford-Bohr model

5. Heisenburg Uncertainty principle

6. Wave mechanical atom

7. Schrodinger “Quantum model”

8. Orbital theory

9. Electronic configuration theory

The 9 atomic structure

Thomson model "Raisin Bread Model''

• atoms is a spherical mass containing electrons and that this spherical mass is positive but is made neutral by the electrons embedded in it

• -negatively charged particles are embedded in positively charged particles

Rutherford model

• it is based on additional experimental evidence of “alpha scattering experiments”

• Ernest Rutherford used the Geiger-Marsden experiment to prove that most of an atom is in a very small space called the atomic nucleus.

• Rutherford took a photo plate and surrounded it with gold foil, and then shot alpha particles at it. "Gold Film/Foil Experiment"

Bohr model of an atom

• in this model, protons are in the nucleus and the electrons are in the orbital motion around the nucleus

• Niels Bohr introduced the Bohr model. This model showed that electrons orbit the nucleus in fixed circular orbits. -Planetary Model

Rutherford-Bohr model

in this model, the atoms are in elliptical orbits of increasing number

Heisenburg Uncertainty principle

this principle states that simultaneous determination of the exact position and exact momentum of electron is impossible

Wave Mechanical atom

in this model, the nucleus is a single cluster of particles at the center of the atom while the electrons are everywhere

Schrodinger “Quantum model”

this theory makes the assertion that electromagnetic radiation like X-rays, gamma rays, radio waves and light rays are made up of small bits of energy

Pauli’s exclusion principle

Quantum numbers:

• Principal Quantum Number (n)

• Azithmuthal Quantum Number (l)

• Magnetic Quantum Number (m)

• Spin Quantum Number (s)

Orbital theory

• this theory states that the number of orbital types in a given shell is equal to the shell number

• Orbitals have a three dimensional region in space where the probability of finding the electron is greatest

• Hund's Rule of Maximum Multiplicity

Electronic Configuration theory

1. The 1st main energy level can have a maximum number of 2 electrons.

2. The 2nd main energy level can have a maximum number 8 electrons (2s, 6p)

3. The 3rd main energy level can have a maximum number of 18 electrons (2s, 6p, 10d)

4. The 4th main energy level can have a maximum of 32 electrons (2s, 6p, 10d, 14f)

5. An s-orbital can only have 2 electrons, a p-orbital 6, a d-orbital 10, and an f-orbital 14

Atomic Number (Z)

• number of protons in the nucleus of an atom of an element

• also the number of electrons in an atom

• this quantity is fundamental to the identity of each element because it is related to the electrical make-up of atom, therefore:

Mass number

• total number of protons and neutrons on the nucleus, therefore:

  mass # = # of protons + # of neutrons

Mass number

Nuclear notation (A)

Neutrons

Nuclear notation (N)

Atomic number

Nuclear notation (Z)

Isotopes

• atoms of the same element with the same atomic number, but different mass numbers

• they have the same number of protons and electrons but different number of neutrons

• many elements exist as two or more stable isotopes, although one isotopes is usually present in greater abundance than another isotopes

Isotones

atoms of different elements having the same number of neutrons

Isobars

atoms of different elements having the same atomic mass

Ions

• it is a charged species, an atom or a molecule, that has lost or gained one or more electrons

• Cation

• Anion

Cation

When an atom or molecule loses electrons, it becomes a positively charged ion called

Anion

When an atom or molecule gains electrons, it becomes a negatively charged ion called

Molecules

it is the smallest indivisible portion of a pure chemical substance that has its unique set of chemical properties, that is, its potential to undergo a certain set of chemical reactions with other substances

Anode and Cathode

2 Electrodes

Anode

electrode where oxidation occurs

Cathode

electrode where reduction occurs

Extrinsic property

• are the physical properties of matter which may vary from time to time

• not the characteristics of the substance itself

• depends on the amount, also called extensive properties

• examples are height, weight, temperature, size, shape, volume

Intrinsic property

• are the properties of matter which are constant

• give the characteristics of the substane its unique identity

• those which do not depend on the amount, also called intensive properties

• examples are boiling point, freezing point, melting point, viscosity, refractive index

Mass

• constant at any place and time

• a measure of the quantity of matter in an object

• when travelled to the moon, the mass of an object will still be the same

• can never be zero

Weight

• varies, depends on the amount of gravity

• refers to the downward pull of the objects towards the center of the earth; the force that gravity exerts on an object

• when travelled to the moon, the weight of an object will only be 1/6 of its weight on earth

• can also be zero

• Elements

• Compounds

2 Pure substances

• Metals

• Non-metals

• Metalloids

3 Elements

Oxygen - 45.5%

Silicon - 27.2%

Aluminum - 8.3%

Iron - 6.2%

Calcium - 4.7%

Magnesium - 2.8%

All others - 5.3%

Percentages of the FF in earth’s crust:

Oxygen -

Silicon -

Aluminum -

Iron -

Calcium -

Magnesium -

All others -

Percentages of the FF in the human body:

Oxygen - 65%

Carbon - 18%

Hydrogen - 10%

Nitrogen - 3%

Calcium - 1.6%

Phosphorus - 1.2%

All others - 1.2%

Percentages of the FF in the human body:

Oxygen -

Carbon -

Hydrogen -

Nitrogen -

Calcium -

Phosphorus -

All others -

• Ionic compound

• Covalent compound

• Metallic compound

• Organic/Inorganic compound

• Acids

• Bases

• Salts

• Oxides

8 Compounds

Organic compounds

• contain carbon, usually bonded to hydrogen.

• ex: DNA, Sugar, Methane, Ethanol

Inorganic Compounds

• usually don't contain carbon it includes

  carbon dioxide and some carbonates,

  cyanides, and carbides.

• ex: Table salt, Hydrochloric acid, Quartz, Carbon Dioxide

Acid

Blue litmus turns red

Base

Red litmus turns blue

Acids

• Taste sour

• React with some metals to give off hydrogen gas

• Conduct electricity in solution

Bases

• Taste bitter

• Feel slippery

• Dissolve fats and oils

The pH of a solution is a measure of H+ ion concentration in a solution. On a scale from 0 to 14, lower pH solutions are strong acids; higher pH solutions are bases. Solutions measuring pH 7 are neutral.

• Homogenous

• Heterogenous

• Solutions

• Colloids

• Suspension

• Emulsions

6 Mixtures

Homogeneous Mixture

• particles distributed uniformly

• ex: Vodka, Steel, Air, Rain

Heterogeneous Mixture

• particles distributed non-uniformly

• ex: Cereal in milk, Ice in soda, Soil, Blood

True solution

Particle size less than 10^-7 cm

Colloidal solution

Particle size between 10^-7 cm and 10^-5 cm

Suspensions

Particle size Greater than 10^-5 cm

Tyndall effect

the scattering of light by particles in a colloid or suspension

Not visible

Is light source visible in this mixture?

Solution (water)

Visible

Is light source visible in this mixture?

Colloid (milk)

Visible

Is light source visible in this mixture?

Suspension (Flour and water)

Compounds

• always have a definite composition by weight

• Preparation shows evidence of chemical action taking place

• components can be separated by chemical means

• constituents can be separated by chemical means

• composed of two or more substances that are chemically combined

Mixtures

• components may be present in any proportions

• it is prepared with no evidence of any chemical reaction taking place

• components do not lose identity

• components may be separated by mechanical means

• composed of two or more substances that are not chemically combined

Law of Definite proportions

a chemical compound always contains exactly the same proportion of elements by mass

Law of Multiple proportions

• when two elements combine to form more than one compound, the weights of one element that combine with a fixed weight of the other are in a ration of small whole numbers

Law of Combining weights

• elements combine in the ratio of their combining weights or chemical equivalents; or in some simple multiple or sub-multiple of that ratio.

• also called the Law of Reciprocal proportions or Law of Equivalents