Exam 1: Ch. 2 - Atoms, molecules, and Ions

Dalton’s Atomic Theory

• Elements are composed of indivisible particles called atoms.

• All atoms of the same element are identical (Size, mass, chemical properties, etc.)

• Compounds are composed of different elements.

• Chemical reactions involve the separation and combination of elements. But elements are not created or destroyed.

-Some of these points of the theory are outdated.

Law of Multiple Proportions

The same elements can be part of a compound in different ratios. This difference in ratios makes compounds different.

ex) CO vs CO2, H2C2 vs CH4 vs C2H6, H20 vs H202

*Compounds look similar but are very different

Law of Conservation of Mass

• Mass is not created or destroyed. The same number of atoms exist when reactants react to become products. As on the example below, we have CH4 reacting with O2 to form H2O and CO2.

• We have the same number of carbon, oxygen, and hydrogen atoms (one, four, and four, respectively), on the reactant side (left side of the arrow) as well as on the product side (right side of the arrow).

CH4 + O2 + O2 → H2O + H2O + CO2

Components of the Atom

• Atoms are made up of subatomic particles.

• These are protons, neutrons, and electrons.

• Protons are positively charged, neutrons have no charge, and electrons are negatively charged.

• The mass of the proton is about the same as the neutron (1.67 x 10-24 g), but they are both regarded as 1 amu (amu = atomic mass unit) when dealing with an atom’s mass.

• The mass of an electron is 9.0 x 10-28 g. Since it is MUCH smaller than the proton and neutron, the electron’s mass does not contribute to an atom’s mass.

Atomic Number

The number of protons that an atom has is what identifies the atom.

Isotopes

Atoms of the same elements (same number of protons) but contain a different mass.

This is due to these elements having a different number of neutrons.

Atoms, molecules, and compounds

What are the components of matter?

Atoms

The basic unit of matter.

Molecules

Atoms bonded together through a chemical bond.

Compounds

Made up of different elements.

Pure

• Not combined with anything else.

100% H2O

100% CO2

Heterogeneous

Composition varies

Homogeneous

Composition is exactly the same.

Proton

1 amu

Positive (+) charge

Neutron

1 amu

No charge

Electron

No mass

Negative (-) charge

protons + neutrons

Mass = _________________

Column 1

Alkali Metals (1 electron)

Column 2

Alkaline Earth Metals (lose 2 electrons and form a +2 charge)

Column 16

Elements pick up 2 electrons, stable with a -2 charge.

Column 17

Halogens (can pick up one electron to form a -1 charge)

Column 18

Noble Gases (they don’t react with anything, most stable)

Columns 3-12

Transition Metals

Ion

Atom/molecule with a charge.

Cation

Atom/molecule with a positive charge.

Anion

Atom/molecule with a negative charge.

Ionic Compounds

• These compounds are composed of metals and non-metals.

• Metals carry the positive charge, the non-metals carry the negative charge.

ex) Na + Cl → Na+ + Cl-

ex) Ca + 2F → Ca2+ + 2F-

• Ionic compounds do not form discrete units, but rather an array of alternating positive and negative charges.

Smaller

As neutral atoms lose electrons, their radius becomes _____________.

Larger

As neutral atoms gain electrons, their radius becomes ____________.

NOTES

Cations and anions can exist as atoms or molecules.

Molecules and Molecular Compounds

• Two or more atoms connected through chemical bonds.

• Chemical bonds are made up of two electrons each bond, atoms can form single, double and triple bonds.

• Non-metals form these types of bonds with other non-metals.

• Molecules are discrete units that together give the bulk substance its properties.

Ex) H-H, H-O-H, N≡N, O=C=O

Molecule

Two or more atoms covalently bonded together. The atoms can be the same element or different elements.

Ex) O2 → two oxygen atoms bonded together.

Ex) H2O → two hydrogen atoms and one oxygen atom.

Molecular compound

A substance made of two or more different nonmetals that are covalently bonded together.

Ex) H2O → hydrogen + oxygen

Ex) CO2 → carbon + oxygen

NOTES

Ionic Formulas

• When writing an ionic formula, the cation always goes first.

• The name of the cation stays the same.

• The name of the anion changes by adding -ide at the end.

o Chlorine becomes chloride, fluorine becomes fluoride, oxygen becomes oxide, nitrogen becomes nitride, etc.

• Since we mostly focus on the metals to the far left of the periodic table and the non-metals, we will skip the transition metal properties. Therefore, the numbers at the top of the columns change to the bolded numbers at the top of this table.

Examples of ionic formulas:

Al2O3, CaBr2, MgO, Cs3N, BaS

Writing a formula from ions:

Mg2+ + N3- → Mg3N

Ti4+ + O2- → TiO2

Carbonate

Permanganate

Ammonium

Sulfate

Phosphate

Nitrate

Hydroxide

Sodium carbonate

Calcium nitrate

Strontium hydroxide

Potassium sulfate

Barium hydroxide

Rubidium phosphate

NOTES

Ionic Compounds with Transition Metals

• Transition metals can form multiple charges that are stable

o Fe2+ , Fe3+ , Cu+ , Cu2+ , etc.

• Older nomenclature:

ex) Fe2+ , Ferrous

ex) Fe3+ , Ferric

ex) Cu+ , Cuprous

ex) Cu2+ ,Cupric

ex) FeCl2 Ferrous chloride

ex) FeCl3 Ferric chloride

ex) Cu2O Cuprous oxide

ex) CuO Cupric oxide

• This older method is not very consistent. NO3- Nitrate OH- Potassium sulfate Hydroxide Barium hydroxide Rubidium phosphate

• Updated nomenclature uses the metal’s name (unchanged) with Roman numerals to denote the charge on the metal.

ex) FeCl2 Iron(II) chloride

ex) FeCl3 Iron(III) chloride

ex) Cu2O Copper(I) oxide

ex) CuO Copper(II) oxide

Lead (II) hydroxide

Iron (II) nitrate

Lead (II) oxide

NOTES

• Use prefixes when there is more than one of the same atom

ex) Di-, tri-, tetra-, penta-, etc.

• The larger number of atoms goes last (PBr5 Vs. BrF3)

1

Mono

2

Di

3

Tri

4

Tetra

5

Penta

6

Hexa

7

Hepta

8

Octa

9

Nona

10

Deca

Bronsted Acid

Donates a proton (H+)

Bronsted Base

Accepts a proton (H+)

NOTES

Molecular acids of halides

• These acids dissociate in water, water takes the H+ from the acid.

• The name of the acid before it dissociates is hydrogen chloride.

• When it dissociates in water, the name changes to hydrochloric acid.

ex) H-Cl → H+ + Cl-

ex) H-Br → H+ + Br-

ex) H-I → H+ + I-

NOTES

Oxoacids

• These acids contain H, O, along with a different element at the center of the structure.

• The formulas are written with H first, followed by the central atom, and oxygen last.

Common oxoacids:

H2CO3

H3PO4

HClO3

H2SO4

Empirical Formulas

• The simplest form of a chemical formula.

Examples:

The molecular formula for caffeine is: C8H10N4O2

Caffeine’s empirical formula is: C4H5N2O

The molecular formula for hydrogen peroxide is: H2O2

Hydrogen peroxide’s empirical formula is: HO

The molecular formula for dimethyl fumarate is: C8H12O4

Dimethyl fumarate’s empirical formula is: C2H3O

NOTES

The roman numeral for naming a compound is only the positive charge.

Number of electrons

____________ = number of protons - charge

NOTES

The Roman numeral corresponds to the charge of the cation.

NOTES

For the purposes of this exam, we are going to say that only ionic compounds (metal and nonmetal) can be simplified to its lowest form.

So the compound H2O2 CANNOT be simplified to HO because this is covalent (nonmetal and nonmetal)

Ions are charged particles formed by the loss or gain of electrons, while molecules are neutral groups of two or more atoms bonded together.

What is the difference between ions and molecules?