Interphase Chemistry - Week 1

Mixture

Can be separated into components by physical means (melting,freezing, boiling, dissolving)

Homogenous mixture

Mixture where composition is uniform throughout (air, steel, saltwater)

Heterogeneous mixture

Mixture where components aren’t uniformly distributed (oil/water, milk, etc)

Law of conservation of mass

The mass of a compounds formed = sum of masses of component elements

Law of definite proportions

A chemical law stating that a given compound always contains its component elements in fixed ratio by mass, regardless of the sample size or source. This means that the relative proportions of the elements in a compound are constant, leading to predictable mass relationships in chemical reactions.

Which of the postulates from Dalton’s atomic theory of matter still holds

true with our modern knowledge of chemistry?

1. Matter consists of indivisible atoms:

NO – nucleus, electrons, protons, quarks...

2. All the atoms of a given chemical element are identical in mass and in all other

properties:

NO - Isotopes

1. Different elements have different kinds of atoms; their atoms have different

masses:YES

2. Atoms are indestructible and retain their identities in chemical reactions:

NO – nuclear reactions

3. A compound forms from its elements through the combination of atoms of unlike

elements in small whole-number ratios:

NO – non-stoichiometric compounds

How to find the empirical formula given the percentage of masses for each compound

Divide percentages of masses of each element present in compound,

Law of combining volumes

When gases combine in a chemical reaction at fixed temperature and

pressure, the volumes of the reacting gases and products are in

simple integer ratios.

• While mass is conserved, volume is not

• 1 L of oxygen + 2 L of hydrogen → 2 L of water vapor

• Since we know that atoms combine in simple integer ratios to form

molecules, and we see these same integer ratios when gases

combine, the simplest explanation is that they are the same ratio!

Mass Number (A)

The total number of protons and neutrons in an atom's nucleus, which determines its isotopic identity.

Atomic Number (Z)

The number of protons in the nucleus of an atom, which determines the element's identity and its position in the periodic table.

Average Atomic Mass

The weighted average of the masses of all the isotopes of an element, measured in atomic mass units (amu). It reflects both the mass and relative abundance of each isotope. (sum of percentage times the mass of each)

Avogadro’s Number

The number of particles in one mole of a substance, approximately equal to 6.022 x 10²³. It is used to convert between moles and individual particles, such as atoms or molecules. The number of atoms in exactly 12 g of 12C

Molar Mass

1 mol of any element weighs the average atomic mass in

grams

• Molar mass - the mass of one mole of particles (atoms or

molecules) of an element or compound

• Units of grams per mole (g mol-1)

• Calculated by summing the dimensionless relative atomic masses

of the component elements multiplied by their quantity in the

molecular formula

Rounding sig figs

When the number after 5 is even, round down bcs you cant get it up

Addition and Subtraction sig figs

Round answer to the least precise amount of decimal places (includes, ones, tens, hundreds places)

Multiplication and division sig figs

Least number of sig figs

Logs and exponents sig figs

Answer to log as the same number of places as decimal as original number had in sig figs. reverse for exponents (5.032² = 25.3 (3 decimal places, 3 sig figs)

Relative Atomic Mass

The dimensionless mass of an atom  measure on relative

scale with the mass of 12C equaling 12.

• Ex: Relative atomic mass of 23Na is: 22.98977. What does it

mean?

• 23Na is 22.98977/12 times as heavy as 12C

Empirical Formula Vs Molecular Formula

Empirical formula: the simplest formula that gives correct and

simplified ratio of atoms of each kind in a compound

• Useful for solids and ionic compounds, helpful in combustion analysis

• Molecular formula: specifies # of atoms of each element in one

molecule

Chemical Formula and percent composition

Example: Ethylene (C2H4)

• Empirical formula: CH2

Its composition by mass can be calculated based on 1 mol of the compound:

Mass of C = (1 mol C)(12.011 g/mol) = 12.011 g

Mass of H = (2 mol H)(1.000794 g/mol) = 2.0159 g

Total mass of 1 mol of compound is 14.027 g

• Each mass % is then found by dividing each of the elemental masses by the molar

mass and multiplying by 100 %

Final answer: Ethylene is 85.625 % C and 14.372 % H

4 basic types of of reactions

4 Basic Types of Reaction:

1. Synthesis – A + B  AB

A compound made from simple materials

2. Decomposition – AB  A + B

A compound broken down into smaller compounds/elements

3. Combustion – CxHy(Oz) + O2  xCO2 + yH2O

A compound containing carbon & hydrogen (sometimes oxygen). Combines w/ O2 (g) to

produce CO2 + H2O.

4. Replacement

*Usually Redox rxns  later in course.

To find formulas of compounds given percentages

D1) Divide mass percentage of one by the other, for all compounds

2) Divide each ratio by the smallest ratio

3) this will give the ratio of each

1. Assume a 100-gram Sample: This allows you to directly convert the percentages into grams.

2. Convert Percentages to Grams: If the compound is 20% carbon and 80% oxygen, you would have 20 grams of carbon and 80 grams of oxygen.

3. Convert Grams to Moles: Divide the mass of each element by its molar mass to find the number of moles.

4. Determine the Mole Ratio: Divide the number of moles of each element by the smallest number of moles calculated. This will give you the ratio of atoms in the empirical formula.

5. Write the Empirical Formula: If the mole ratios are not whole numbers, multiply them by a factor to get whole numbers. These whole numbers represent the subscripts in the empirical formula.

6. Calculate Empirical Formula Mass: Add up the atomic masses of all the atoms in the empirical formula.