Chapter 1-3 Summer Work

Matter

• Solid has a fixed shape and volume

• Liquid has a fixed volume but takes the shape of the container

• Gas has neither fixed volume nor shape

Pure substances

• Element or compound

Element

• Cannot be broken down into two or more pure substances ( it’s already in it’s purest form)

Compounds

• A pure substance that contains more than one element

• Fixed composition ( always contains the same elements in the same mass percentages —> ex water always has 2 hydrogen molecules and 1 oxygen molecule)

• Properties of compounds are different than its elements that make it up —> ex: Sodium Chloride is unreactive but sodium is extremely reactive and chlorine is poisonous

• techniques such as electrolysis and heat are used to separate the compound into its pure substance

• Uniform properties and lack of physical mixing with other substances

Mixtures

• Contain two or more substances that retains chemical identity ( physically combined but not chemically) so they contain their original chemical properties

Homogenous Mixtures

• When the composition of the mixture is the same throughout. Another name is a solution ( a solvent and a solute)

Examples: salt water, vinegar, coffee without milk, soft drinks

Heterogenous mixtures

• When the composition of the mixture varies throughout such as rocks ( can contain different % of mass of elements/ substances)

• Differing components/ non– uniform

How to separate components of a mixture: Filtration

• Filtration separates a HETEROGENEOUS solid – liquid or solid– gas mixtures. It uses filer paper where the liquid part goes through and the solid stays in the filter paper.

How to separate components of a mixture: Distillation

• Distillation separates HOMOGENOUS solid–liquid or liquid– liquid mixture. The liquid vaporizes, leaving a residue of solid in the distilling flask ( uses heat)

• Heats the mixture to vaporize the more volatile ( lower boiling point) which travels to a condenser. Then it is cooled down back to a liquid.

How to separate components of a mixture: Chromatography

• Chromatography separates HOMOGENOUS Gas liquid or solid liquid mixtures

• Gas Chromatography vaporizes a liquid or solid as it travels through the column

• Compounds with lower boiling point move through the column faster, resulting in a chromatogram, a graph that plots intensity versus time.

Units of Measurement and instruments

• Standard Unit of length of is Meter.

• Volume is measure in cubic centimeters, liters, and milliliters.

• mass is expressed in grams, kilograms, or milligrams.

Mass

• A measure of the amount of matter in an object

Weight

• A measure of the gravitational force acting on the object

Temperature

determines the direction of heat flow

• Heat flows higher temp to lower temp

• SI Unit is kelvin

Significant Figures

• Used when there is uncertainty of the measurements when measuring volume

• Meaningful digit obtained in a measurement

  • RULES for Sig Figs: —> Non– Zeros, Zeros between digits, and zeros after the digit with decimal points are significant ( zeros in front of digits aren’t)

Rules for adding and subtracting, and multiplying and division

• For multiplying and dividing the number of sig figs is the same as the smallest number of sig figs in a number

• For adding and subtracting the number of decimal places which is the smallest is the amount of sig figs you should have ( ex .355 + .8976 would lead the total to have 3 sig figsa

Rules for rounding off

• If the digit is to be discarded is less than 500 leave it alone

• if the digit to be discarded is more than 500 round up

• If the digit to be discarded is = 500 round so its an even number

Conversion of units

1m=100 cm

1m=1000mm

1m=1 × 10^9 nm

1 ft =0.305 m 12 in

1 yd =0.914 m 3 ft

1 mi=5280 ft

1 in =2.54 cm

1 cm³= 1 ml, 1 l =1.0567 qt, 1 ft³= 28.32 L

1 kg=1000 g

1 g=1000 mg

1 lb= 16 oz

1 lb=453.6g

Properties of substances:

• Intensive: Must be independent ( separate) of the amount of substance and doesn’t change even different mass such as density, boiling point, etc.

• Extensive: Properties that depend on the amount of substance such as mass and volume

Chemical Properties:

• Properties that occur when the substance takes part in a chemical reaction where it is changed to a new substance.

• Examples are: Flammability, toxicity, reactivity with other substances, heat of combustion, and the ability to break down other substances ( Corrosivity)

Physical Properties:

• Properties without changing the chemical identity of a substance.

• Examples are: Melting point —> the temperature at which substances change from liquid to solid ( Substances don’t change identity when melted just their states change)

• Boiling point —→ the temperature at which vapors ( a gas forms)

• Density

Units of conversion

1 Kilometer =1000 meters

1 meter= 10 decimeters

1 meter=100 centimeters

1 meter=1000 milimeters

1 meter=10^6 micrometers (u)

1 meter=10^9 nanometers (n)

Atoms:

An Element composed of tiny particles. All atoms of a given element, have the same chemical properties, while atoms of different elements have different properties.

Dalton Theory

• The Atom can be defined as the smallest particle of an element that can enter a chemical reaction

• He believed All atoms of a given element are identical in mass and properties

• In ordinary chemical reactions atoms move from substance to substance but no atom of any element disappears or is changed into another atom.

• In a compound the relative number of atoms are integers or simple fraction.

• Different combinations produce different compounds

• Atoms of different elements have different masses

Components of the atom

• JJ Thomson discovered the Electron in which particles passed through the cathode ray tube. The rays or particles deflected from the positive pole showing that it is negatively charged.

• Electrons carry a -1 charge and weight 1/2000.

The Law of conservation of mass

• John Dalton discovered that mass can’t be created or destroyed.

Law of constant composition

• Compounds always contain the same elements in the same proportions by mass, regardless of how prepared.

Law of multiple proportions

• John Dalton discovered that Two elements combine to form different compounds, the ratios of the masses of one element to another are whole numbers

Protons and Neutrons

• JJ Thomson suggested the plum pudding model where there is a sphere of positive charge and negatively charged electrons embedded like plums on top of pudding.

• Rutherford did an experiment where he passed/ fired a particles ( helium atoms without their electrons) and most went through the foil without changing their direction but a few reflected back at acute angles.

• Scattering was caused by a positively charged nucleus at the center of the foil’s atom. Most of the a atoms is empty space which is why the bouncing occurred vs passed through

• The proton has a mass of nearly 1 carries a +1 charge. The neutron an uncharged particle with mass also a bit more than 1.

Atomic number

• All atoms of a particular element have the same number of protons —> symbol z = # of protons

• in a neutral atom the # of protons equals the # of electrons

Atomic masses

• How heavy atoms of different elements are

Mass numbers; isotopes

• Mass number is represented by A= # of protons + # of neutrons

• All atoms of a particular element may differ in mass because the # of neutrons can be different

• Istopes: Atoms with the same # of protons but different # of neutrons

• To Find the average atomic mass you do mass number of the element ( times the percentage of abundance/100) + the other elements.

Isotopic Abundances:

• The parentage of a specific isotope within a naturally occurring element —> total adds up to 100% —> used to calculate the element’s average atomic mass.

• By comparing the accelerating voltages required to bring ions to the same point on the detector, it is possible to determine relative masses of ions.

Mass Spectrometry:

• The atomic masses ( relative abundance) of the two isotopes are determined by the proportions of the heights of the peaks —> the x– axis determine that the isotopes % have the quantity of mass ex) Cl —> 75% have a mass of 35 amu

Masses of individual atoms; avogardo’s number

Na= 6.022 × 10²³ ( represents the number of atoms of an element in a sample whose mass in grams is numerically equal to the atomic mass of the element)

Periods and Groups

• The horizontal rows are periods ( ex: Li-Ne)

• The vertical column, 1 - 18 are groups

• Groups 3–12 are transition metals

• Groups 1,2,13,14,15,16,17,18 are main group elements

• Group 1 is alkali metals

• Group 2 us alkaline earth metals

• Group 17 is Halogens

• Group 18 is noble gases

  *The periodic table is an arrangement of elements, in order of increasing atomic number. Elements in vertical groups have similar chemical properties.

• Dimitri Mendeleev invented the periodic table

Metals and nonmetals

• Groups 1 – 16, except H, B, Is, Ge are metals

• Metals have high electrical conductivities

• Boron, Silicon, Ge, Arsenic, Sb, Te are metalloids because their properties fall between

• Groups 14– 18 are non– metals

Molecules

• Two or more atoms combine to form an uncharged molecules

• Within the two atoms ( commonly as non– metallic forces) called covalent bonds are sharing electrons —> intermolecular forces ( the ones that hold separate molecules together are quite weak)

Structural formulas

• Shows the bonding pattern within the molecule

Condensed Structural formula

• Suggests the bonding pattern in the molecule and highlights

the presence if a reactive group of atoms within the molecule

Ions

• When an atom loses or gains electrons, giving it a net charge

• Metal atoms lose electrons and form cations

• Non metals gain electrons and form anions

• The atomic number doesn’t change when an ion is formed

• Polyatomic ions are held by covalent bonds

Ionic bonds

• Strong electrical forces holding oppositely charged ions ( a cation and an anion)

• Ionic compounds are solids and have high relatively melting points.

• To melt ionic compounds they must be separated —> When dissolved they conduct an electric current making it a stronger electrolyte

Discrete molecules

• Separate units, held be strong covalent bonds, neutral

• attracted to each other by intermolecular forces

• Aren’t solids, rather move around

Ocet rule

• Atoms that are close to a noble gas in the periodic table form ions that contain the same number of electrons as the neighboring noble gas atom to try to form a stable electron configuration.

• transition and post– transition metals form positive ions and more than one.

Polyatomic ions positive

• Ammonium ( NH₄⁺⁾

• Hydronium ( H₃O⁺⁾

Polyatomic -1

• Acetate ( C₂H₃O₂^-1)

• Hydroxide (OH^-1)

• Nitrate ( NO₃^-1)

• Chlorate ( ClO₃^-1)

• Nitrie ( NO₂^-1)

polyatomic -2

• Carbonate ( CO₃^-2)

• Sulfate(SO₄^-2)

polyatomic -3

• Phosphate (PO₄^-3)

Binary Molecular Compound

• Product of when two nonmetals combine with each other

• Uses the prefixes di,tri,tetra,penta,hexa,hepta,octave,nona,and deca

• The second half ends it wide

Naming ionic compounds

• When a nonmetal forms two oxanions ate is used for the anion with the larger number and its is used for the anion containing few oxygen atoms.

Naming acids

• If name of anion ends in ide then the acid name is hydro___ ic —> Ex)HCL —> hydrochloric acid

• If name ends in ite(polyatomic) than acid name is ___ ous

  —> EX) HNO₂ —> nitrous acid

• If name ends in ate than acid name is ____ic

  —→ EX) H₂SO3 —→ Sulfurous acid

• Acids are substances that ionize ( gain or lose an electron) in water to produce hydrogen ions (H⁺)

Percent composition

• Found by the mass of the element present and the mass of the substance

  —> EX) 11.19 g H/100 = 11.19% H

• When given mass percents in a compound assume 100 grams

Combustion

Mass of sample= mass of O+ mass of C + mass of h

OR mass of O= mass of sample - ( mass of C + mass of H)

Theoretical yield

• The maximum amount of product that can be formed in a chemical reaction

• Found by using the limiting reactant and whatever you get as the total for the product is theoretical yield

Percent yield

• actual yield ( experimental)/theoritcal x 100

How to find excess reactant’s mass

• Can find it by subtracting the amount of moles in the product to the amount of moles of the excess reactant and converting

  OR by using the mass of the limiting reactant converting to moles of the excess and multiplying by the molar mass to find the amount of mass used up. Subtract that by the mass given in the problem.

Mass Spectra of diatomic elements

Elements that are diatomic such as Chlorine, oxygen, nitrogen, etc. may have more lines on their mass spectra because their atoms may stick together