Chemistry 111 Vocab Flashcards Chapters 1 + 2

Chemistry

The science that seeks to understand the behavior of matter by studying the behaviors of atoms and molecules

Macroscopic

Things we can observe

Microscopic

Things we cannot observe but can visualize

Symbolic

Special language that represents both domains

Mass

a measure of the amount of matter in an object

weight

(N) a measure of g-force on an object

Law of Conservation of Mass

Mass is neither created nor destroyed in ordinary chemical and physical changes

Pure substances have constant composition

True, and they are elements and comounds

Elements

cannot be broken into simpler substances by chem

1 type of element

Compounds

can be broken down

2+ types of elements

Mixture

is composed of two or more types of matter that can be present in varying amounts; homogenous and heterogeneuous

Homogenous mixture

uniform composition

Heterogeneous mixture

Composition varies

Molecules

2+ atoms connected by a chemical bond

Physical property

can be observed without changing the composition of a substance

Physical change

is a change in the state or properties of matter without any accompanying change in its chemical composition

Chemical property

a property that a substance displays only by changing its composition via a chemical change

Chemical change

the molecular composition changes; a reaction!

Extensive Properties

Depends on the amount

can be added up

Examples include:

• volume

• mass

• size

• weight

Intensive Properties

Does not depend on the amount

cannot be added up

Examples include:

• Boiling Point

• Color

• Temperature

• Luster

• Hardness

Density

Mass/Volume

All the digits in a measurement, including the uncertain last digit, are called

Significant Figures

Significant Figure Rules

Determining significant figures:

. All nonzero digits (1-9) are always significant: 28.03; 0.0540

2. Interior zeros are significant: 408; 7.0301

3. Leading zeros are not significant: 0.0032; 0.00006

4. Trailing zeros ➢ Significant if they are after the decimal place: 45.000; 3.5600 ➢Significant if they are before a decimal place and after a nonzero digit: 1050.;140.00

5. Exact numbers have an infinite number of sig figs.

6. Every digit in scientific notation is significant: 1.12 x 103

Multiplication and division in Sig Figs

The product is reported with the same number of sig figs as the factor with the fewest sig figs

Addition and subtraction in Sig Figs

The product carries the same number of decimal places as the quantity with the fewer decimal places

Rounding in Sig Figs

Round down if the last digit dropped is 4 or less: 5.34 ~ 5.3

Round up if the last digit dropped is 5 or more: 5.37 ~ 5.4

A measurement is said to be ____ if it yields very similar results when repeated in the same manner

A measurement is considered _________ if it yields a result that is very close to the true or accepted value

Precise

Accurate

Dimensional analysis

the mathematical approach of using units as a guide to solving problems

Units are subjected to the same mathematical operations as their associated numbers

Conversion factor

a ratio of two equivalent quantities expressed with different measurement units

Fahrenheit and Celsius

Tf = (9F/5C * Tc) + 32

Kelvin and Celsius

Tk = Tc +273.15

Atomic Mass Unit (amu)

1 amu = 1.6605 × 10^-24 g

1/12 of the mass of one carbon 12 atom

Fundamental Unit of Charge (e)

e = - 1.602 × 10–19 C

Atomic number (Z)

# of protons in an element. Defines the element!

Mass number (A)

# of protons + # of neutrons

A - Z = # of neutrons

A neutral atom has # of protons = # electrons

so Z = # of electrons (e-)

Isotope

same element (Z) but different number of neutrons

Atomic mass

average mass of all naturally occurring isotopes of an element

Molecular formula shows: 1) ______ to indicate the types of atoms 2) ______________ after the symbol to indicate the number of each atom

Chemical symboks

Subscripts

Structural formula

shows how atoms are connected

Elements that exist as molecules

Diatomic molecules: H2, N2, O2, F2, U2, Br2, I2, S8

The Mole

mol = 6.022 × 10²³ particles

aka Avogadro’s number

Molar mass

The mass in grams of 1 mole of the substance (g/mol)

The molar mass of a compound is the mass in grams of 1 mol of its molecules

Formula Mass

The sum of the average atomic masses of all the atoms in the molecular formula

Covalent substances exist as discrete molecules

Formula mass can be referred to as a molecular mass

Formula Mass for Ionic Compounds

• Ionic substances are composed of discrete cations and anions combined in ratios to yield electrically neutral compounds

• Ionic compounds do not exist as molecules

• Formula mass cannot be referred to as a molecular mass

• The average atomic masses of the ions can be approximated to be the same as the average atomic masses of the neutral atoms

Percent Composition

first step is often to measure its constituent elements in the process of determining the formula experimentally

the percent by mass of each element in a compound

Mass % = (total mass of element / mass of compound) * 100%

Determination of Empirical Formulas

1st way: Grams of atoms in compound * molar mass → Moles of atoms in compound → Divide by smallest mole amount to determine the whole number mole ratio

2nd way: Assume you have 100.0g of a compound → conert each percentage to grams → Determine moles of each element after dividing by mol mass → Determine whole number mole ratio after dividing by the smallest mole amount

Derivation of Molecular Formulas

A compound’s molecular formula can be determined from its empirical formula and its molecular or molar mass

molecular mass/empirical formula mass = n formula units/molecule

______: another term used for a homogeneous mixture – uniform composition and properties throughout its entire volume

• The relative amount of a given solution component is known as its _____

• A solution consists of two components:

• _____: component with a greater concentration

• ______: component at a much lower concentration than the solvent • A solution in which water is the solvent is called an _____ solution

Solutions

concentration

Solvent

Solute

aqueous

Concentration of a solution

the amount of solute present in a given amount of solvent

Molarity

(M) – the number of moles of solute in exactly 1 liter (1 L) of the solution:

M = mol solute/ L solution which is n/V

Dilution of Solutions

the procedure for preparing a less concentrated solution from a more concentrated solution

Concentrated: relatively high concentration

Dilute: relatively low concentration

Mass percentage

the ratio of the component’s mass to the solution’s mass, expressed as a percentage

mass percentage = mass of component/mass of solution * 100

Also called: percent mass, percent weight, weight/weight percent • Symbols: %, %mass, %weight, (w/w)%

Volume Percentage

The concentration of a solution formed by dissolving a liquid solute in a liquid solvent is often expressed as that

Volume % = volume solute/ volume solution * 100%

Mass - Volume Percentage

A ratio of a solute’s mass to the solution’s volume expressed as a percentage

%m/v

The specific units used for solute mass and solution volume may vary, depending on the solution

Parts per million

ppm = mass solute/ mass of solution * 10^6

Parts per billion

ppb = mass solute / mass of solution * 10^9