Chemistry lecture study for test 1

Mass

the amount of matter in an object (how much stuff is in the object)

Matter

a substance or solution that occupies space and has mass

Weight

The force gravity exerts on an object

Law of Conservation of Matter

Matter cannot be created nor destroyed

(Matter will have the same mass even when it has changed form)

Chemical changes

atoms recombine to form new substances

Physical substance

Matter that changes states of matter but not its composition

Elements

Pure substances that cannot be broken down into simpler forms by chemical changes

Compounds

Pure substances that can be broken down

Consists of 2+ elements that’s chemically bonded

Compound properties are different from element properties

Law of constant composition

if you take a small piece of an element, it will still have the properties of the entire element

Mixtures

Composed of 2 or more types of matter that can be present in varying amounts and can be separated in varying amounts

Homogenous (solutions) mixtures

Mixtures that are so uniform you can’t tell the matters apart

Ex. Air, coffee, brass

Heterogenous mixtures

Easy to tell one matter from another and easy to separate

Ex. oil and water, sand, smog

Pure substance can’t be…

Mixtures

Atoms

Smallest part of an element that has the properties of the element and can enter chemical combination

Molecules

2 or more atoms connected by strong forces called chemical bonds

Compounds are…

molecules but not all molecules are compounds

Ex. H₂ is a molecule but not a compound because it is composed of only one element

Properties

characteristics that help distinguish one substance from another

Physical properties

Physical changes

Ex. density, color, durability, melting/boiling points and electrical durability

Chemical properties

Chemical changes, can’t change state of matter

Ex. Flammability, toxicity, acidity, reactivity, and heat of combustion

New substance after reaction

Ex. nail rusting, adding sugar to iced tea

Extensive property

Depends on the amount of matter present

Ex. mass, volume, heat

Intensive property

Independent to the amount of matter present

Ex. density, temp

Measurements provide 3 kinds of info

1) A number

2) A unit

3) An indication of the uncertainty of the measurement

SI units

Length-meter (m)

Mass-Kilogram (Kg)

Time-Second (s)

Temp-Celsius (C)

Electricity- Ampere (A)

Amount of substance- Mole (mol)

Luminous intensity-Candela (Cd)

Fractional SI units

Femto (f) 10^-15

Pico (p) 10^-12

Nano (n) 10^-9

Micro (μ) 10^-6

Milli (m) 10^-3

Centi (c) 10^-2

Deci (d) 10^-1

Kilo (k) 10^3

Mega (M) 10^6

Giga (G) 10^9

Tetra (T) 10^14

Volume

How much space is occupied by an object

Volume formula

V = m/d

Density

How heavy something is relative to its size

Density formula

D = m/v

Mass formula

m = density x volume

What is the density of water

1 gram per mL

Uncertain numbers

Quantities derived from measurements other than counting

Significant numbers

• Non-zero digits

• Captive zeros

• Trailing zeros (when right of decimal place)

Non-zero digit

Any number other than zero

Ex. 1,2,3,4,56,64…

Captive zeros

All zeros between non-zero digits

Ex. 73.04, 500007, 9.000000006

Trailing zeros

Zeros to the right of the decimal place with no digits after them

Ex. 9.00050, 0.0000000900000

Non significant numbers

Zeros that don’t add significance to the size of the number. Usually left of the decimal point

Non significant numbers example

Leading zeros or trailing zeros with no present decimal point

Ex. 0.009, 0067, 99000

Calculation rules

1) round results to the same decimal place for +/-

2) Round results to the same amount of digits as the # with the least amount of sig. figs. for x/÷

3) for results with the last digit ≤4 round down, if ≥6 round up. If its 5 round up or down but try and make it an even number

Preciseness

When the measurement is done repeatedly and yields similar results each time

Accurate numbers

Yields results that are close to the true/accepted value

Dimensional analysis

Used to convert one unit to another unit

Conversion factor

A number used to switch units by x/÷

Dalton’s atomic theory (4 points)

1) All elements are composed of atoms

2) All atoms of the same element are identical; different elements have different types of atoms

3) Atoms can’t be subdivided, created or destroyed

4) Atoms of different elements can combine in simple whole number ratios to form chemical compounds

Law of Multiple Proportion

We can form different substances using the same elements just in different amounts

Ex. A green solid has 0.558g Cl to 1g Cu but a brown solid has 1.116g Cl to 1g Cu

Electrons

• First subatomic particle to be discovered

• Negatively charged, attracted to positive charges

• Charge to mass ratio= 1.759 × 10¹¹ C/Kg

• Occupy almost all of an atom’s volume

• Discovered by J.J Thompson using cathode rays

Nucleus

• Small and dense

• Contains most of the atoms mass

• surrounded by electrons

• The center of the atom

• Discovered by Ernest Rutherford

Neutrons

• Uncharged

• Same mass as protons, heavier than electrons

• Found in the nucleus

• Discovered by James Chadwick

Protons

• Positively charged

• Located in the nucleus

• Heavier than electrons

• Discovered by Ernest Rutherford

Isotopes

• Same element but different mass

  • Ex. Carbon 12 vs carbon 13

• Caused by differing levels of protons

• Written by writing the mass number as a superscript to the top left of the element symbol

Atomic number (z)

The number of protons in the nucleus

• this value determines the identity of the atom

• Any atom with 6 protons is carbon regardless of the isotopes

• Also indicates the amount of electrons in a neutral atom

Neutral atoms

An atom with an equal number of protons, electrons, and neutrons

Mass number (A)

Total number of protons and neutrons in an atom

• The number of neutrons is equal to the mass number - the number of protons

Ions

The imbalance between protons and electrons

• Atom charge = # of protons - # of electrons

• Atoms and molecules acquire charge by losing or gaining electrons

Anion

A negatively charged ion

Cation

A positively charged ion

Atomic mass

Each proton and neutron has a mass of ~1 amu

• Electrons weigh far less

The atomic mass of one atom is roughly equal to its mass number

Mass spectrometry

The occurrence and natural abundance of isotopes can be experimentally determined using a mass spectrometer

Molecular formula

A representation of a a molecule or compound and has

1) Chemical symbols

2) Subscripts after the symbol to indicate each atom quantity in the molecule

Ex. CH₄

Structural formula

shows the same things as a molecular formula plus how the molecules connect

Empirical formula

The simplest way of writing IONIC COMPOUNDS (Always in whole numbers)

Molecular formula

Depicts the actual number of atoms in an element in the compound

MF= C₆H₆ Empirical formula= CH

MF= C₂H₄O₂ Empirical formula= CH₂O

Isomers

The same atom with different structures

Ex. Acetate Acid and Methyl Formate both have the same Molecular formula

Periodic law

When elements are arranged in order of increasing atomic numbers, there is periodic repetition of their chemical and physical properties

Periods or series

Horizontal rows 1-7

Groups

Vertical columns numbered 1-18

Metals

Shiny, malleable, good conductors of heat and electricity

Metalloids (semi-metals)

Conduct heat and electricity somewhat well and have some properties of metas and some of nonmetals

Nonmetals

Appear dull and are poor conductors of heat and electricity

Monoatomic ions

Only have one atom

Polyatomic ions

multiple atoms with a charge

Oxyanions

Polyatomic ions that contain an oxygen

-ate suffix means

there’s more oxygen

-ite suffix means

there’s less oxygen

Per- prefix means

Largest amount of oxygen (more than -ate)

Hypo- prefix means

Smallest amount of oxygen (less than -ite)

Ionic bonds

Transfer of electrons (losing or gaining electrons)

• Non contact forces

Covalent bonds

Happens when electrons are shared and molecules form

• Contact force

Ionic compounds

• Metals become positive and form cations

• Non metals lose electrons and form anions

• Metals and nonmetals form ionic compounds

• Held together by ionic bonds

• Occurs a lot in transition metals or main group metals combining with a non metal

• Solid with high melting and boiling points

• Non-conductive in solid form

• Conductive in molten form

Formulas of ionic compounds

• Write the symbol and charge of the metal first and the non metal second

• Swap the charges (They will be even)

  • Ex. Li^1 O^-2 → Li₂ O

• Reduce to the lowest ratio and write the new charges on bottom right of the elements

• Many ionic compounds contain polyatomic ions as the cation, anion or both

• Parenthesis in ionic compound formulas mean there are two or more poly atomic ions

Ex. Ca²⁺ and PO₄^3- forms Ca₃(PO₄)₂

Naming Ionic compounds with a metal ion with a variable charge

• Most of the transition metals and some main group metals can form a 2+ cations with different charges

• The charges of the metal ion is specified by a roman numeral in parenthesis after the name of the metal

  • roman numerals indicate positive charge

• Older naming systems used the -ous and -ic suffixes

  • -ous was for the lower charge

  • -ic was for the higher charge

Ex. FeCl₂ = Iron (II) Chloride