Chapter 2- Atoms

HW-2.10, 2.11, 2.15, 2.17, 2.25, 2.29, 2.45, 2.48, 2.53, 2.54, 2.66,

2.79

Ancient Thinkers

• Pondered what matter was made of• Democritus and followers believed matter

was composed of tiny particles– Called particles atoms– Said their were different types of atoms that

add different properties

• Zeno of Elea believed matter was infinitely divisible

• Atoms- the basic unit of matter

Classifications of Matter

MatterAnything that occupies space and has mass

Pure SubstancesFixed composition; cannot

Be further purified

MixturesA combination of two

Or more pure substances

ElementsCannot be

Subdivided byChemical or

Physical means

CompoundsElements united

In fixed ratios

HomogeneousMatter uniformComposition throughout

HeterogeneousMatter nonuniform

composition

Definitons

• Elements- a substance that consists of identical atoms

• Compounds- a pure substance made up of two or more elements in a fixed ratio by mass.

• Formula- gives the ratio of a compounds constituent elements and identifies each element by its atomic symbol

• Mixture- a combination of two or more pure substances

Types of Mixtures

• Homogeneous mixture- composition the same throughout

• Example: Air – 78% Nitrogen 21% Oxygen

• Heterogeneous mixture- nonuniform composition

• Example: Blood

• Note: Mixtures consists of two or more pure substances with different physical properties. This properties can be used to separate the mixture.

Dalton’s Atomic Theory

1. All matter is made up of very tiny, indivisible particles called atoms

2. All atoms of the same element have the same chemical properties

3. Compounds are formed by the chemical combination of two or more different kinds of atoms

4. A molecule is a tightly bound combination of two or more atoms that acts as a single unit

Laws

• Law of Conservation of Mass– Matter can be neither created nor destroyed

(This does not contradict Dalton!!)

• Law of Constant Composition– Compounds are always made up of elements

in the same proportion by mass

Types of Elements

• Monoatomic Elements- elements that consist of single atoms not connected to each other. Ex. He, Ne

• Diatomic Elements- elements that exist as pairs of atoms, bonded together.

Ex. O2 Cl2 Br2 I2 H2 N2

• Polyatomic Elements- elements that exist with more than two atoms bonded together. Ex. O3 P4 S8 Diamond

Subatomic Particles

• Proton- subatomic particle found in the nucleus of an atom and has a positive charge

• Electron- subatomic particle found in the space surrounding the nucleus and has a negative charge

• Neutron- subatomic particle found in the nucleus and has no charge

Atom Properties

• Mass Number- the sum of the number of protons plus the number of neutrons

• Atomic Number- the number of protons in the nucleus of an element

• Symbol for an atomic nucleus- The element symbol with the mass number as a superscript on the left and the atomic number as a subscript on the left.

Properties of Atoms (cont.)

• Isotopes- atoms with the same number of protons but different numbers of neutrons– Properties of isotopes of the same element

are almost identical except for radioactivity properties.

• Atomic weight- of an element in the Periodic Table is a weighted average of the masses (in amu) of its isotopes found on Earth.

Periodic Table

• First created by Mendeleev• Still used today• Started by arranging elements by weight

beginning with hydrogen• Noticed trends so he started arranging them in

rows (periods) and started a new period every time he found an element with properties like hydrogen.

• Discovered that elements in other groups had similar properties as well.

Definitions

• Main-group elements- elements in groups 1A, 2A, and 3A-8A.

• Transition elements- elements in B groups

• Inner transition elements- elements 58-71 and 90-103

Classes of Elements

• There are three classes

• Metals- most elements in the Table. They are solid (except Hg), shiny, conductors, ductile (wires), and malleable (sheets)

• Alloys- solutions of one or more metals dissolved in another metal. Examples: Bronze (copper and tin) Pewter (tin, antimony, and lead)

Classes (cont)

• Nonmetals- with the exception of Hydrogen, they are on the right side of the table.

• Non conductive (exception graphite)

• Tend to accept electrons in reactions

• Metalloids- B, Si, germanium, arsenic, antimony, and tellurium

• Have some properties of both

Trends in the Periodic Table• Sizes increases as you move down a

column• For the halogens, group 7A, MP and BP

increase as you move top to bottom• Ionization energy- is a measure of how

difficult it is to remove an electron from an atom in the gaseous state

• Ionization energy increases as you move up a group and from left to right across a period

Electron Configuration

• Electrons are located in the considerably larger space outside the nucleus.

• The lowest possible energy level an electron can be in is the Ground State.

• The energy of electrons is quantized and there are only certain energy levels an electron can be in.

Electron Configuration

• Electrons in atoms are confined to specific regions of space called Principal energy levels or more simply, shells.

• Shells are numbered 1,2,3,4, and so on

• Electrons in the first shell are closest to the nucleus, are held most strongly, and therefore are the lowest in energy.

• As you move away, energy increases

Electron Configuration

• Shells are divided into subshells designated by the letters s, p, d, and f.

• Within subshells electrons are grouped into orbitals.

• Orbital- is a region of space that can hold two electrons.

• The first shell has one orbital called the 1s• The second shell contains one s orbital

and three p orbitals.

Electron Configuration

• P orbitals come in sets of three and holds a total of six electrons

• The 3rd shell contains one s orbital, three p orbitals, and five d orbitals.

Orbital Shapes

• All s orbitals have the shape of a sphere with the nucleus in the middle.

• Each 2p orbital has the shape of a dumbbell with the nucleus at the midpoint of the dumbell.

• The three 2p orbitals are at right angles to each other.

Electron Configuration

• Electron Configuration- a description of the orbitals that its electrons occupy

• In the ground state of an atom only the lowest energy orbitals are occupied, all others are empty

• We determine the ground state configuration using the following rules:

Rules• Orbitals fill in the order of increasing

energy from lowest to highest

• Each orbital can hold up to two electrons with spins paired

• When there is a set of orbitals of equal energy, each orbital becomes half-filled before any of them becomes completely filled.

3 types of configuration notation

• Orbital box diagrams- use boxes or lines to represent an orbital

• Condensed configuration- write only shell number, orbital, and number of electrons in that orbital

• Noble Gas Notation- use noble gas symbol for inner shell designation

Showing electron config.

• Usually only show the outer shell electrons• Outer-shell electrons are called valence

electrons• The energy level in which valence

electrons are found is called the valence shell.

• To show the outmost electrons of an atom, we commonly use a representation called a Lewis Dot Structure

Lewis Dot Structures

• Lewis Dot Structures- shows the symbol of the element surrounded by a number of dots equal to the number of electrons in the outer shell of an atom of that element.

• Valence electrons- are electrons in the outermost shell.

• Examples-

Another trend in the Periodic Table

• Elements in a group have the same electron configuration in their outer shell!!

• That mean the Lewis Structure will be the same as well, except for the atomic symbol.

Atoms vs. Ions• Atoms are neutral species that have the

same number of electrons and protons.

• During typical chemical reactions, the number of protons and neutrons does not change

• The number of electrons does change!! Atoms can gain or lose electrons.

• Ions are a compound or element that has gained or lost an electron and is now charged!!

Atoms vs. Ions• Example-

• Ionization Energy- a measure of how difficult it is to remove an electron from an atom.

• Ionization Energy increases as you move up a group and from left to right on the periodic table.