Cycle 6 Chemistry

French Scientist Made a list of all 33 known elements at that time Broke into 4 categories

Gases Metals Non-Metals Earths

Arranged elements by increasing mass and putting elements that shared properties in columns Realized that there were undiscovered periodic elements Considered the Father of the Periodic Table

Arranges Elements by Atomic Number Created the Periodic Law

Elements in the same vertical column Otherwise known as families Numbered 1-18

Elements in the same horizontal row Numbered 1-7

Elements in groups 1,2, and 13-18 (The Main Points of Reference for what we are doing)

States that there is a periodic repetition of physical and chemical properties of the elements when arranged according to increasing atomic number

Increasing atomic number

Defined as electrons in the outermost orbitals of the atom

Shiny Solid at Room Temperature Good Conductors of heat Most elements are metals

Elements that are generally gases or brittle Poor conductors

Have physical and chemical properties of both elements 8 Metals

Alkaline Metals Very Reactive

Alkaline Earth Metals Very Reactive

Similar Properties to other metals but do act a bit differently due to their partially filled subshells

The two columns at the bottom of the periodic table Lanthanides Actinides

Halogens Highly reactive Very Close to becoming Noble Gases

Noble Gasses Unreactive In a perfect state of balance

By the element at the top of the group Ex. Carbon Family

S, D, P, F

Atoms of an element which gained or lost electrons

Atoms that gain electrons and become negatively charged. More electrons than protons Atoms get bigger as they gain more electrons

Atoms that lose electrons and become positively charged. Atoms get smaller as they give away their valence electron. More protons than Electrons

To help increase stability of the atoms

The size of the atom is influenced by the number of electrons that an atom has. Atomic size is determined by how close an atom is to its neighboring atom. Definition: Half the distance between two equidistance nuclei

The atomic radius increases as you go to the left in the periodic table and down.

Larger Orbital Size More Electron shielding

Indicates the relative ability of its atoms to attract electrons in a chemical bond

As you move more to the right and up.

Electronegativity increases as you move to the right More protons in the nucleus to attract electrons

Electronegativity increases as you go up Strong Attraction to electrons

Atoms tend to gain, lose, or share electrons in order to achieve stability. Aka 8 valence electrons

This is because the Periodic Table is periodic, meaning there is a repeating pattern with it.

IONIC BONDS FORM BETWEEN METALS AND NON-METALS Ionic Bonds from when atoms become ions

Formula Units

The chemical formula of anionic compounds that lists the ions in the lowest ratio that equals a neutral electric charge.

NOOOOOOOOOOOO

In a Crystal Lattice Structure

A repeating pattern of positive and negatively charged ions

Maximizes attraction between positive and negative ions which equal strong bonds Presence of particular metal ions lead to different colors

LATTICE ENERGY DUE TO THE CRYSTAL LATTICE STRUCTURE

Relates the size and charge of the ions to the energy level needed to make the bond

Relates attraction to size and charge. Smaller atoms can be closer together and form stronger bonds. Larger atoms form smaller bonds that are weaker

Do not conduct electricity Charged ions can't flow and make a current

When dissolved in water, they become electrolytes Conduct electricity very well Ions are separated and mixed

Melting Point- Temperature that solids turn to liquids Ionic compounds have a very high melting point because of strong ionic bonds Boiling Point- Temperature at which a liquid becomes a gas. Ionic Compounds have a high boiling point

Hard- Ionic Compounds in the solid form are hard because of strong ionic bonds that hold them together Brittle- Ionic Solids, despite being hard, are brittle of impacted with enough force.

Bonds are broken or made energy is either absorbed or released

Energy is absorbed by the chemical reaction

Energy is released by the chemical reaction THE FORMATION OF IONIC BONDS ARE ALWAYS EXOTHERMIC

TRASH

Ions made with more than one atom bonded together These ions are charged molecules Act as a group

Metals are not ionically bonded but share several properties with ionic compounds. Metals are arranged in crystal lattice structures in their solid states. The attraction between the positive metal cations and the negatively charged sea of electrons

The primary explanation for metallic bonding. Metals are described as having Localized metal cations (positively charged) Delocalized valence electrons (negatively charged)

All metals contribute to their valence electrons to form a sea of electrons. The delocalized electrons are free to move around the metal cations.

Related to the number of valence electrons contributed by the metal atom. The more valence electrons, the stronger the metallic bond.

Metals are able to be hammered or rolled into thin sheets Ex. Aluminum foil

Metals are able to be drawn into thin wires

Mixtures of metals that have unique properties.

Alloys capable of taking the place of other metals because of their similar properties

Intertwined within the solids. In the smaller places electrons would reside in, but not as small as an electron.

The last number in an electron configuration will match up to whatever PERIOD an element is in.

Defined by how closely an atom lies to a neighboring atom ATOMIC SIZE AND ATOMIC RADIUS ARE DIFFERNENT

They become bigger

Involving the shielding of electrons and their charge from the nucleus of an atom.

Can pull electrons closer to the nucleus of an atom. The stronger it is, the more electrons can be pulled.

The smallest arrangement of atoms in a crystal lattice structure that has the same symmetry as the whole crystal; a small representative that is part of a whole