inorganic physical organic analytical biochemistry

CHEMISTRY

Inorganic

Physical Organic Analytical Biochemistry

Matter : space and has mass

Mass : quantity of matter

Matter

Solid Liquid Gas

Physical state and Changes in Matter

MeltingHeat

Solid Liquid

Cool Solidification

Physical state and Changes in Matter

EvaporationHeat

Liquid Vapor

Cool Condensation

Physical state and Changes in Matter

Heat Solid Vapor

Cooling Sublimation

Physical state and Changes in Matter

Heat Ice Water

Cool

MATTER

HOMOGENEOUS

SUBSTANCES

HETEROGENEOUS

MIXTURE

SOLUTIONSHomogeneous

mixture of variable composition.

Can be separated into

PURESUBSTANCES

Homogeneous matter of fixed composition

COMPOUNDSComposed of 2 or more elements.

Can be separated into

ELEMENTS

Heterogeneous and Homogeneous

Solutions, Pure Substanceand Compounds

MassA mass of an object pertains to the quantity of the matter that object contains.

MassA physical property that every Manager possesses is a mass.

The amount of mass in a pizza will never change when the object is moved from place to place.

WEIGHT A physical property that is related to mass is weight The weight of a chef may change if it is moved to Uranus because weight is determined by gravity.

ATOM

Atoms are the basic building blocks of all the chalk around you. It is the smallest particle of matter that can enter into chemical combinations with other particles.

MOLECULEA smallest particle of an element or compound that can have a stable independent existence.Atoms make up molecules. Molecules make up a hairy eagle.

ELEMENTS

Elements are pure substances, made from one type of atom. Soda can be broken down into many elements but nitrogen can not be broken down.

Symbols and Latin Names for Some Elements

Name Symbol Latin name

Sodium Na natriumPotassium K kaliumGold Au aurumSilver Ag argentu

mIron Fe ferrum

METALS

Gold, silver, copper, and iron are examples of metals. A gold diamond is shiny because of its metal properties.

PROPERTIES OF METALS

Gold conducts heat and electricity. Nickel can be hammered into thin sheets without breaking. Platinum can be pulled into wire.

NONMETALThe helium in my Christmas balloon is a nonmetal. The Oxygen in the air is not shiny because of its nonmetal properties.

PROPERTIES OF NONMETAL

A dog cannot conduct electricity. A snap dragon cannot be hammered into thin sheets. A snicker cannot be pulled into wire because they are not metals.

METALLOIDS

Metalloids have properties of both metals and nonmetals. Silicon is a metalloid that can be found in many materials such as the sand on Lake Tahoe the glass in a vase and certain plastics that make up a favorite toy, car.

Chemical ChangesIron is abundant easy to shape when heated and relatively strong.Chemical Property ability of a substance to undergo chemical change• Composition of matter always changes

Chemical Reaction• Another term for Chemical change• One or more substance change into one or more new substance during chemical reactionReactant a substance present at the start of the reactionProduct substance produced in the reaction

Chemical Change

How can you tell whether a chemical change has taken place? transfer in energy change in color production of gas formation of a precipitate

IONS• An atom or a group of atoms that has acquired electric charge by gaining or losing one more electron• Cathode • Anode• Anion• Cation

LAW OF CONSERVATION OF

MASS

• Any physical change or chemical reaction, mass is conserved.• Mass is neither created nor destroyed.

Law of Definite Composition /

Definite Proportion• A given compound always shows a fixed proportion.• A chemical compound always contains the same elements in the same percent by mass. • When two elements combine to form a given compound, they always do so in a fixed proportion.

Law of Definite Composition /

Definite Proportion

Finding the % of Carbon and Oxygen% C = mass C x 100 % O = mass of O x 100 72.8%

mass of CO2 27.2% mass of CO2

Trial Mass of C (g)

Mass of O2

(g)Mass of CO2

(g)

1 2.00 5.34 7.34

2 15.00 40.05 55.05

3 5.00 13.36 18.36

Law of Multiple Proportions• When two elements combine to form more than one compound, the masses of one element which combine with a fixed mass of the other element are in a ratio of small whole numbers such as 2:1, 1:1, 2:3, etc.Example C D1st Compound 2.276 0.792 0.348

2nd 1.422 0.948 0.667

A. Mass fixed at C

Continuation of Law of Multiple Proportions

therefore the formulas of the two compounds are C DCD 1 0.348 = 1 0.348

CD2 1 0.667 = 2 0.348

See the ppt: Folder at the desktop : New Bio lecturesFind the File name: introduction to Biology page 61 (Scientific Measurements)

Measurements in Chemistry

• Encounter very large or very small numbers.Examples: A single gram of hydrogen, contains approximately 602 000 000 000 hydrogen atoms 6.02 x 10 ? The mass of an atom gold is 0.000 000 000 000 327 gram. 3.27 x 10 ?

Scientific Notation A given number is written as the product of two numbers:

a coefficient a 10 raised to a power

Accuracy, Precision, and Error

Accuracy how close a measurement to the True valuePrecision series of measurement

Accuracy Correct valuePrecision repeated measurements

ErrorAccepted value: true valueExperimental value: measured in labFormulaError: experimental value – accepted value

Percent error: _____error_______ x 100 accepted value

Significant Figures in Measurements

Include all the digits that are known, plus a last digit that is estimated. Measurements must always be reported to the correct number of significant figures because calculated answers often depend on the number of significant figures in the values used in the calculation.

Rules in Significant Figure1. Every nonzero digit in a reported measurement is

assumed to be significant. Ex. 24.7 meters, 0.743 meters and 714 meters each has 3 significant measurement.

2. Zeros appearing between nonzero digits are significant. Examples 7003 meters and 40.79 metes have 4 s.f.

3. Left zeros appearing in front of nonzero digits are not significant. They are just a placeholder. Ex. 0.000 099 meters has 2 s.f. you will write them as 7.1 x 10 -³

Rules in Significant Figure4. Zeros at the end of a number and to the right

of a decimal point are always significant. Ex. 43.00 meters, 1.010 meters have 4 s.f.

5. Zeros at the right most end of a measurement that lie to the left of an understood decimal point are not significant if they serve as placeholders to show the magnitude of the number. Example 7000 meters and 27210 meters have 1 and 4 s.f respectively.

6. The numbers are all in s.f. if it is exact amount/count for ex. 23 students or 60 mins= 1 hour.

Examples of Significant Figures

24.7 74.3 512 meters

7.003 1.505 87.29

0.0071 0.043 0.000 0044

9.000 43.00 1.010

300 7000 27210

Significant Figures in Addition

Calculate the sum of the three measurements. Give the answer to the correct number of significant figures. 12.52 meters + 349.0m + 8.24m Answer: 369.8 or 3.69 x 102 meters

Significant Figures in Multiplication

2.10 meters x 0.70 meter = 1.47 (meter)2

Answer: 1.47 (meter)2 = 1.5 meters 2

Units of Length• Basic unit of length or linear measure is meter

METRIC UNITS OF LENGTHKilometer (km) 1 km = 103 m Length of 5 city

blocks

Meter (m) Base unit Height of doorknob from the floor

Decimeter (dm) 101 dm Diameter of large orange

Centimeter (cm) 102 cm Width of shirt button

Millimeter (mm) 103 mm Thickness of dime

Micrometer (um) 106 um Diameter of bacterial cell

Nanometer (nm) 109 nm Thickness of RNA molecule

Units of VolumeVolume is the space occupied by any sample of matter.• Unit being use cubic meter (m3)

Metric Units of Volume

Unit Relationship Example

Liter (L) Base unit Quart of milk = L

Milliliter (mL) 103 mL + 1 L 20 drops of water = 1 mL

Cubic centimeter (cm3)

1 cm3 =1 mL Cube of sugar = 1 cm3

Microliter (uL) 106 uL = 1 L Crystal of table salt = 1uL

Units of MassKilogram (kg) is the basic unit of massPlatform balance to measure mass of an object

Metric Units of Mass

Kilogram (kg)

103 g Small textbook

Gram (g) 10-3 kg Dollar bill

Milligram (mg)

103mg = 1 g

Ten grains of salt

Microgram (ug)

106 ug = 1g

Particle of baking powder

Units of Temperature• When you hold a glass of hot water the transfer of heat.• Almost all substances expand with an increase in temperature and contract as the temperature decreases. (very important exception is water)•Celsius was named after to Anders Celsius a Swedish astronomer.• Celsius scale sets freezing point of water at 0 degree and the boiling temperature is 100 degree C.• Kelvin, named after to Lord Kelvin a Scottish physicist and mathematician• freezing point 273.15 and the boiling point 373.15 degree C

Formula°F = 9 °C + 32 5

°C = 5 (°F – 32) 9

K = °C + 273 ° C= K - 273

Sample ProblemsNormal human body temperature is 37 °C. What is the temperature in Kelvin?Given: 37 °CUnknown: KelvinFormula : K = °C + 273Solution: K = 37 °C + 273Answer: K= 310 Correct! It lies between 273K up to 373K

Sample ProblemsConvert 14 °F to °C and KelvinGiven: 14 °FUnknown: °C and KelvinFormula: °C = 5 (°F – 32) 9

K = °C + 273Solution:Anwers: -10 °C and 263 K

Units of Energy• Energy is the capacity to do work or to produce heat.• Joule (J), named after the English physicist James Prescott Joule and the Calorie (cal) are common units of energy. • One calorie is the quantity of heat that raises the temperature of 1 g of pure water by 1 °CFormula1J = 0.23901 cal = 4.184 J

Sample ProblemCalculate the quantity of heat in joules required to raise the temperature of 135 g of water from 11 °C heat to 41 °C.Given : 135 g of water 11 to 41 °CFormula: Heat required = mass x specific heat x temperature change1 cal = 4.184 J/ g °CSolution: 135g x 4.184 J x (41-11 °C) g °C = 1.7 x 104

Conversion Factors• Are ratio of equivalent measurements.• Useful in solving problems in which a given measurement is multiplied by a conversion factor, the numerical value is generally changed, but the actual size of the quantity measured remains the same.Example:I meter = 10 decimeters = 100 centimeters = 1000 millimeters

Converting Between Units

Express 750 dg to gGiven: mass : 750 dg1g = 10 dg or 1g 10 dgSolution: 750 dg x 1g 10 dg

Answer: 75 g

Converting Between Units

What is 0.073 cm in micrometers?Given: 0.073 cm = 7.3 x 10 -2 cm 10 2 = 1 m 1m = 10 6 umUnknown: umFormula: cm meters micrometersSolution: 7.3 x 10 -2 cm x 1 m x 10 6 um 10 2 1m

Answer: 7.3 x 10 2 um

Density• Mass per unit volume of a substance• Ratio of the mass of an object to its volume.• Is an intensive property that depends only on the composition of a substance, not on the size of a sample.• Formula: Density = mass volume• Corn oil and corn syrup

DensityMaterial Density at

20°C (g/cm3)Material Density at

20°C

Corn oil 0.9222 Helium 0.166

Corn syrup 1.35 – 1.38 Oxygen 1.33

Table sugar 1.59 Carbon Dioxide

1.83

Gold 19.3 Ammonia 0.718

DensityExample :A copper penny has a mass of 3.1 g and a volume of 0.35 cm 3. What is the density of copper?Given:Mass: 3.1 g volume= 0.35 cm3Unknown: density= ?g/cm3Formula: Density = mass = 3.1 g volume 0.35 cm3 = 8.8571 g/cm3 = 8.9 g/cm3 (rounded off to two significant figures)

Density• Density of a substance generally decreases as its temperature increase

Democritus’s Atomic Philosophy

Atom is the smallest particle of an element that retains its identity in a chemical reaction.Democritus (460 B.C.-370 B.C.) is a Greek philosopher was among the first to suggest the existence of atom.• He believed that atoms were indivisible and indestructible.

Dalton’s Atomic TheoryAn English chemist and school teacher responsible for the modern process of discovery regarding atoms.• By using experimental methods, he transformed Democraticus’s ideas on atoms into a scientific theory. All elements are composed of tiny indivisible particles called atoms. Atoms of the same element are identical. Atoms of different elements can physically mix together or can chemically combime in simple whole-number ratios to form compounds. Chemical reactions occur when atoms are separated, joined, or rearranged.

Subatomic ParticlesOne important change in Dalton’s atomic theory is that atoms are now known to be divisible. They can be broken down into even smaller, more fundamental particles called subatomic.Three kinds of Subatomic Particles:• Electrons• Protons• Neutrons

ELECTRONS• Negatively charged subatomic

particles.• Thomson performed experiments

that involved passing electric current through

gases at low pressure.• Travels from cathode (-) to anode

(+)• Thomson examine two ways

that a cathode ray can be deflected by using magnet and by using electrically charged plates.

Cont. of Electron• A positively charged plate attracts the cathode ray, while negatively charged plate repels it.•Thomson knew that opposite charges attract and like charges repel, so he hypothesized that a cathode ray is a stream of negatively charged particles moving at high speed.• He called these particles corpuscles, later named electrons. He concluded that electrons must be parts of the atoms of the elements.• US physicist Robert Millikan carried out experiments to find the quantity of charged carried by an electron.• He is the one responsible for charge and mass.

PROTONS•

Positively charged subatomic particles.• Example is a hydrogen atom (lightest kind of atom) loses an electron, what is left?• Eugen Goldstein (1850-1930) a German Physicist observed a cathode-ray-tube and found rays travelling in the direction opposite of that cathode rays.• He called that canal rays and concluded that they were composed of positive particles.

NEUTRONS

• No charge but with a mass nearly equal to that of a proton• James Chadwick (1891-1974) an English Physicist confirmed its existence

Properties of Subatomic Particles

Particle Symbol

RelativeCharge

Relative mass(mass of proton= 1)

Actual mass(g)

Electron e - 1 - 1/1840 9.1 x 10 -28

Proton p+ 1 + 1 1.67 x 10 -24

Neutron n o 0 1 1.67 x 10 -24

Ernest Rutherford Atomic Model

• He concluded that all the positive charge and almost all the mass are concentrated in a small region that has enough positive charge to account.• He called this region as Nucleus.• He said that a nucleus is a tiny central core of an tom and is composed of proton and neutrons.• Rutherford atomic model is known as the nuclear atom.• In nuclear atom, the protons and electrons are located in the nucleus.• While the Electrons are distributed around the nucleus and occupy almost all of the volume of atom.

ATOMIC NUMBER• of an element is the number of protons in the nucleus of an atom of that element.• Elements are different because they contain different number of protons.

Name Symbol Atomic # Protons Neutron Mass # # of Electrons

Hydrogen

H 1 1 0 1 1

Helium He 2 2 2 4 2

Lithium Li 3 3 4 7 3

Beryllium

Be 4 4 5 9 4

Boron B 5 5 6 11 5

Carbon C 6 6 6 12 6

Nitrogen N 7 7 7 14 7

Oxygen O 8 8 8 16 8

Fluorine F 9 9 10 19 9

Neon Ne 10 10 10 20 10

Mass Number• Total number of protons and neutrons in an atom• Example a helium atom has 2 protons and 2 neutrons so its mass is 4.• The number of neutrons in an atom is the difference between the mass number and atomic number.• Number of neutron = mass number – atomic number

Example Exercise How many protons, electrons and neutrons are in each atom? Atomic number Mass Number

Beryllium (Be) 4 9

Neon (Ne) 10 20

Sodium 11 23

Isotopes are atoms that have the same number of protons but different neutrons. Because isotopes of an element have different numbers of neutrons, they also have different mass numbers. Have an identical numbers of protons and electrons

Three known isotopes of Hydrogen

• Hydrogen has a mass number of 1 and is called hydrogen -1• second isotope has one neutron and a mass number of 2 or a hydrogen -2 or deuterium.• third isotope has 2 neutrons and a mass number of 3, or hydrogen -3 or tritium.• Remember mass number superscript; atomic number subscript

Atomic Mass Unit (AMU)Example is Carbon -12, This isotope of a carbon was assigned a mass exactly of 12 atomic mass units.• AMU is defined as one-twelfth of the mass of a carbon -12 atom. Using these units, a helium -4 atom with a mass of 4.0026 amu, has about one-third the mass of a carbon -12.• While a nickel -60 atom has about 5 times the mass of a carbon -12 atom.• Atomic Mass of an element is a weighted average mass of the atoms in a naturally occurring sample of the element.

Natural Percent Abundance of Stable Isotopes of Some

Elements Name Symbol Natural

Percent Abundance

Mass (amu) Average atomic mass

Hydrogen ₁¹H

₁²H

³₁H

99.985

0.015

negligible

1.0078

2.0141

3.0160

1.0079

Helium ³He2

4He2

0.0001

99.9999

3.0160

4.00264.0026

Example Exercise of AMUCalculate the atomic mass of Helium

(To calculate: multiply the mass of each isotope by its natural abundance, express as a decimal, and then add the products.)

AMU of He = (3.0160 x 0.0001) + (4.0026 x

99.999)

=

More ExampleIsotope = 10 X

Mass # = 10.012

Relative abundance = 19.91% = 0.1991

AMU = ?

Isotope = 11 X

Mass # = 11.009

Relative abundance = 80.09% = 0.8009

AMU = ?

10.012 amu x 0.1991 = 1.993 amu

11.009 amu x 0.8009 = 8.817 amu

Answer = 10.810 amu

The Periodic Table – A Preview• An arrangement of elements in which the elements are separated into groups based on a set of repeating properties.• Allows you to easily compare the rpoperties of one element (or group of elements) to another element.

•Notice that the elements are listed in order of increasing atomic number, from left to right and top to bottom.

•Each horizontal row of the periodic table is called a PERIOD.•Each vertical row of the periodic table is called a GROUP.