Chem 4.1-4.3 objectives

naming

how do ionic compounds form?

• happens between a metal + nonmetal.

• Electrons are transferred.

• Metal loses electrons → cation (+).

• Nonmetal gains electrons → anion (−).

• Opposite charges attract and form an ionic compound.

cation

positive ion

anion

negative ion

example of ionic bond

• Na → Na⁺ + e⁻

• Cl + e⁻ → Cl⁻

• Na⁺ + Cl⁻ → NaCl

how do covalent bonds form?

• Usually happens between nonmetal + nonmetal.

• Atoms share electrons.

• The shared electrons form a covalent bond.

• The resulting substance is a molecule.

covalent bond example

• H₂O → oxygen shares electrons with two hydrogen atoms

Breaking bonds requires energy. →

Energy goes in.

A bond forms when the overall arrangement becomes

more stable/lower in energy.

Electronegativity

how strongly an atom attracts shared electrons

electronegativity increases when

you move left to right across periodic table

electronegativity decreases when

you move top to bottom across periodic table

Bonding isn't always simply

"ionic" or "covalent."

The difference in electronegativity (ΔEN) helps determine

the type of bond

Nonpolar covalent

• Electrons are shared almost equally.

• Small or zero electronegativity difference.

Polar covalent

• Electrons are shared unequally.

• One atom attracts the electrons more strongly

Ionic

• Electrons are essentially transferred.

• Large electronegativity difference

Polyatomic ion

a group of atoms covalently bonded together that has an overall charge

In a polar covalent bond, the atom with the higher electronegativity has

the higher electron density

The atoms inside the polyatomic ion are connected by _____ bonds, but the entire ion can form an _____ bond with another ion

covalent, ionic

The more electronegative atom gets a

a partial negative charge (δ−)

The less electronegative atom gets

a partial positive charge (δ+)

Acid

Usually starts with H when written as a formula

Before nomenclature you have to

Idientify if it’s ionic,covalent, or acid

Which group on the periodic table can have multiple charges?

transition metal

transition metals

Valence electrons: Varies
Common charges: Multiple positive charges
Characteristics: Transition metals are metals that are generally hard, strong, shiny, and good conductors of heat and electricity. They usually have high melting points and can form multiple ions with different charges. Many transition-metal compounds are colorful. They are generally less reactive than Groups 1 and 2. When naming compounds, a Roman numeral tells you the charge of the transition metal.
Examples: Fe²⁺ = iron(II), Fe³⁺ = iron(III), Cu⁺ = copper(I), Cu²⁺ = copper(II)

alkali metals

Valence electrons: 1
Common charge: +1
Characteristics: Alkali metals are very reactive metals because they have only 1 valence electron, which they easily lose to become more stable. They are soft, have relatively low melting points, and react strongly with water. Their reactivity increases as you move down the group. They are almost never found as pure elements in nature because they react so easily.
Example: Na → Na⁺ + e⁻

alkaline earth metals

Valence electrons: 2
Common charge: +2
Characteristics: These are reactive metals that have 2 valence electrons. They usually lose both electrons to become stable and form +2 ions. They are harder and less reactive than Group 1 metals. Their reactivity increases as you move down the group.
Example: Mg → Mg²⁺ + 2e⁻

boron group

Valence electrons: 3
Common charge: +3
Characteristics: Group 13 elements have 3 valence electrons and can lose those electrons to form +3 ions. Boron is a metalloid, while the elements farther down the group are mostly metals. Metallic character increases as you move down the group. They can form both ionic and covalent compounds.
Example: Al → Al³⁺ + 3e⁻

carbon group

Valence electrons: 4
Common charges: Can vary; often form covalent bonds
Characteristics: Group 14 elements have 4 valence electrons. Because gaining or losing 4 electrons requires a lot of energy, these elements often share electrons and form covalent bonds. Carbon is a nonmetal, silicon and germanium are metalloids, and tin and lead are metals. Metallic character increases as you move down the group.
Important: There is not one simple charge that applies to every Group 14 element.

nitrogen group

Valence electrons: 5
Common charge: −3 for many main-group ionic compounds
Characteristics: Group 15 elements have 5 valence electrons. They need 3 more electrons to complete their outer shell, so they can gain 3 electrons and form −3 ions. They can also share electrons and form covalent bonds. Nitrogen and phosphorus are nonmetals, arsenic and antimony are metalloids, and bismuth is a metal.
Example: N + 3e⁻ → N³⁻

oxygen group/chalcogens

Valence electrons: 6
Common charge: −2
Characteristics: Group 16 elements have 6 valence electrons and usually gain 2 electrons to achieve a full outer shell. They commonly form −2 ions. Oxygen and sulfur are important nonmetals. Metallic character increases as you move down the group.
Example: O + 2e⁻ → O²⁻

halogens group

Valence electrons: 7
Common charge: −1
Characteristics: Halogens are very reactive nonmetals because they only need 1 more electron to complete their valence shell. They usually gain 1 electron and form −1 ions. They react readily with metals to form salts. Halogens are diatomic, meaning they naturally occur as pairs: F₂, Cl₂, Br₂, I₂. Their reactivity decreases as you move down the group. Fluorine is the most reactive halogen.
Example: Cl + e⁻ → Cl⁻

noble gases group

Valence electrons: 8, except He which has 2
Common charge: 0
Characteristics: Noble gases have a full valence shell, making them extremely stable and generally very unreactive. They are gases at room temperature and normally exist as individual atoms, rather than molecules. They do not normally gain or lose electrons, so their common charge is 0.
Examples: He, Ne, Ar

lanthanides

Valence electrons: Varies
Common charge: Usually +3
Characteristics: Lanthanides are metals located in the top row underneath the main periodic table. They are shiny, conductive, and relatively reactive metals. They are often called rare-earth elements and are important in magnets, electronics, batteries, and other technologies.

actinides

Valence electrons: Varies
Common charges: Multiple positive charges
Characteristics: Actinides are metals located in the bottom row underneath the main periodic table. All actinides are radioactive. Many of them are synthetic, meaning they are made artificially. Uranium and plutonium are especially important examples. They can have multiple oxidation states.

Ionic-elements with multiple charges/transition metals naming

(equation to name)

• switch the elements charge

• then use roman numerals to indicate the specific charge

elements with multiple charges/transition metals naming

(equation to name) example

how to know if element has a positive or negative charge

charge=protons-electrons, if the difference is positive it has a positive charge, if it has negative charge it’s negative

how to find amount of protons a element has

looking at the atomic number, they’re equal to each other

how to find amount of electrons a element has

electrons usually match amount of protons unless it has a charge alongside it

covalent naming

• Name the first element.

• Use a greek prefix to symbolize charges, if needed

• use greek prefixes for the second element to symbolize charges, if needed.

• Change the second element's ending to -ide.

polyatomic ions have specific___

names

example of polyatomic ion

NH₄⁺

ammonium

whats NH₄⁺

Ammonium

Ionic hydrate

an ionic compound containing a specific number of water molecules

covalent name example

N₂O₅

→ dinitrogen pentoxide

ionic naming

• write the positive metal or cation first and the negative nonmetal or anion second

• change the nonmetal's ending to "-ide"

ionic naming example

MgO-Magnesium oxide

don’t use “mono” prefix for the _____ element

first

“mono” prefix rule example

NO-Nitrogen Monoxide

1-greek prefix

mono

2-greek prefix

di

3-greek prefix

tri

4-greek prefix

tetra

5-greek prefix

penta

6-greek prefix

hexa

7-greek prefix

hepta

8-greek prefix

octa

9-greek prefix

nona

10-greek prefix

deca

the seven diatomic elements

H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂

the seven diatomic elements mnemonic

Have No Fear Of Ice Cold Beer

• H₂-have

• N₂-no

• F₂-fear

• O₂-of

• I₂=ice

• Cl₂-cold

• Br₂-beer

Binary/Hydro acids naming

Binary acid = H + one nonmetal

Naming:

hydro + root + ic acid

Binary acid name example

HBr → hydrobromic acid

Oxyacid naming

contain H + oxygen-containing polyatomic ion.

Naming rule

-ate → -ic acid

-ite → -ous acid

Oxyacid name example

NO₃⁻ = nitrate
HNO₃ = nitric acid

NO₂⁻ = nitrite
HNO₂ = nitrous acid