introduction to chemistry

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Introduction to Inorganic Chemistry

A Review of the Basic Regents Chemistry Concepts

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Important DefinitionsImportant Definitions


MatterMatter

• Anything that has mass and takes up space

• Everything you “see” around you is composed of matter


MassMass

• The amount of matter an object has

• Measured in grams


ElementElement

• A substance that cannot be broken down chemically into simpler substances

• About 92 naturally occurring elements– Example – gold, magnesium, neon, etc.

• About 25 elements necessary for life• Six most common elements

– Carbon, hydrogen, oxygen, nitrogen . . . calcium and phosphorus

– >99% of living things • Trace elements – necessary in only small (i.e. trace)

amounts– Example – Iodine

• Necessary for proper thyroid function• Goiter caused by iodine deficiency


Goiter Goiter


Normal/GoiterNormal/Goiter


AtomAtom

• Simplest particle of an element that retains the properties of that element

• A given atom is unique to a given element

• Atoms composed of smaller particles called subatomic particles


Compounds vs. MoleculesCompounds vs. Molecules

• Compound - A substance composed of two or more elements.– Ex. NaCl – sodium chloride (table salt)

• Molecule – A substance composed of two or more atoms, can be the same element.– Ex. O2 – oxygen (diatomic)

– However, sodium chloride is also a molecule


The Periodic Table of ElementsThe Periodic Table of Elements((in any language, it’s still the samein any language, it’s still the same))


Important DefinitionsImportant Definitions

• Atomic Number– # of protons– Unique for each element

• Mass number– Sum of the protons and

neutrons

• Atomic Mass– Average of masses of all of

the isotopes of an element– Usually a decimal number

very close to the mass number


Periodic TablePeriodic Table



• Elements arranged according to atomic number starting with hydrogen – (atomic # = 1)



• Contains horizontal rows called periods– Ascending atomic # from

left to right.

• Contains vertical columns referred to as groups.– We are concerned with

groups IA-VIIIA– Elements in a group have

similar chemical and physical properties


Group IA – Alkali MetalsGroup IA – Alkali Metals

• Strong metallic qualities

• Highly reactive– Not found alone in

nature

• One valence electron• Tendency to lose one

e- when reacting.


Group IIA – Alkaline Earth Group IIA – Alkaline Earth MetalsMetals

• Strong metallic qualities

• Very reactive– Not found free in

nature

• Harder and denser than alkali metals

• Two valence e-• Tendency to lose 2 e-

during reactions


Group VIA - ChalcogensGroup VIA - Chalcogens

• More varied in properties– Oxygen, Sulfer –

nonmetals– Selenium, tellurium –

metalloids– Polonium – metal

• Very Reactive • Six valence e-• Tendency to gain two

electrons


Group VIIA - HalogensGroup VIIA - Halogens

• All nonmetals• Often found in

diatomic state (F2)

• Very reactive• Seven valence e-• Tendency to gain 1 e-

during reactions


Group VIIIA- Noble GasesGroup VIIIA- Noble Gases

• Nonmetals• Eight valence e-• Don’t lose or gain

electrons• They are nonreactive

(inert)


Metals vs. NonmetalsMetals vs. Nonmetals


Metals vs. NonmetalsMetals vs. Nonmetals

• Metals– Solid at room temperature– Conduct heat and electricity well– Malleable (sheets)– Ductile (wires)– Lustrous (shiny)– High melting/boiling points

• Nonmetals– Opposite of metals

• Metalloids– Some qualities of both metals and nonmetals


MetalsMetals


NonmetalsNonmetals

• Carbon • Bromine


Trends in the Periodic TableTrends in the Periodic Table


Structure of the AtomStructure of the Atom


Did you know?Did you know?

• Atoms are mostly empty space?– If the nucleus of an atom was the size of a golf

ball, the nearest electron would be roughly 1 km away!

• The nucleus of an atom is extremely dense.– The same size nucleus would have a mass of

approximately 2.5 billion tons!


Some More Things to KnowSome More Things to Know

• All atoms of a given element have the same # of protons

• All atoms are considered neutral in charge unless designated with a symbol of charge, in which case they are considered an ion; # electrons = # protons except in ions

• The # of neutrons is equal to or greater than the number of protons. – mass # - atomic # = # neutrons


3 Major Subatomic Particles3 Major Subatomic Particles

• Proton– Positive charge– 1 x 10 -24 grams (about 1 dalton)– Located in the nucleus

• Neutron– Neutral (no charge)– 1 x 10 -24 grams (about 1 dalton)– Located in the nucleus

• Electron– Negative charge– 1/2000 the mass of a proton or neutron– Moving in orbitals around the nucleus at about the speed of light


Energy Levels (electron shells)Energy Levels (electron shells)

• Electrons exist at varying energy levels• The further they are from the nucleus, the

more energy they have – Think centripetal force

• Electrons tend to occupy the lowest energy level (closest to nucleus) possible

• Electrons can be “excited” to higher energy levels for very brief periods– Example: Light energy during photosynthesis


Excitation of an ElectronExcitation of an Electron


Electron Configuration and Chemical Electron Configuration and Chemical PropertiesProperties

Why atoms reactWhy atoms react• It’s all about the # of valence electrons!

• Valence electron shell is the outermost shell (that contains electrons)

• A full valence shell = inert (stable electron configuration)

• Anything else = reactive (unstable electron configuration)


Here’s the DealHere’s the Deal

• 1st energy level – Full (stable) with 2 electrons

• 2nd energy level– Full (stable) with 8 electrons

• 3rd energy level– Full (stable) with 8 electrons

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Isotopes

• Most elements have at least 2 isotopes, some have several.

• Isotopes vary in the # of neutrons only.

• Example: Carbon has 3 isotopes– 12C – stable (6 neutrons)– 13C – stable (7 neutrons)– 14C – radioactive (8 neutrons)


Uses of Radioactive IsotopesUses of Radioactive Isotopes

• Dating fossils – Carbon – 14

• Measure half-life (5730 years)

• Medical tracers – Iodine – 131

• Various types of sensors can detect radiation.


Shorthand AbbreviationShorthand Abbreviation

• 73Li

• 168O

• How many protons, neutrons and electrons do the above examples have?

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Answers

• Lithium generally has– 3 protons– 3 electrons– 4 neutrons

• Oxygen generally has– 8 protons– 8 electrons– 8 neutrons


Lewis Electron Dot DiagramsLewis Electron Dot Diagrams

• Show the electron configuration for only the valence e- for an atom

• Steps– Write the symbol of the atom– Make a dot for each valence e- (use the “four

sides” of the symbol)– Only one rule – don’t pair up e- until after all

four orbitals have one e- each

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Examples

• Lithium: 1 valence e-

• Chlorine: 7 valence e-

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PracticePractice

• Draw the Lewis dot diagram for the following atoms (use your periodic table)– Hydrogen– Helium– Beryllium– Carbon– Nitrogen– Oxygen– Fluorine– Rubidium– Iodine

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Using LEDD’s to Determine Bonding

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CHEMICAL BONDING


4 Major Types of Bonds4 Major Types of Bonds

• Strongest to weakest– Covalent bonds– Ionic bonds– Hydrogen bonds– van der Waals interactions


Covalent BondsCovalent Bonds• Strongest• Generally occurs when two nonmetals interact• A pair, or pairs, of e- are shared • Single covalent bond

– One pair of e- shared between two atoms – Represented by a single line in structural formula

• Double covalent bond– Two pairs of e- shared between two atoms – Represented by a double line in structural formula

• Triple covalent bond– Three pairs of e- shared between two atoms – Represented by a triple line in structural formula

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Structural vs. Molecular Formulae(single covalent bonds)

• Methane • Methane

CH4

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Double and Triple Covalent Bonds(structural formulae)


Quick PracticeQuick Practice

• React hydrogen with fluorine

• React hydrogen with oxygen

• React hydrogen with carbon

• React carbon with oxygen

• React Nitrogen with hydrogen


Polar vs. Nonpolar Covalent BondsPolar vs. Nonpolar Covalent Bonds

• It’s all about electronegativity– Electronegativity

• The affinity an atom has for electrons– i.e. How strongly it pulls on both its own e- and the e- of other

atoms

• All atoms are electronegative, some more than others

• Polarity, whether or not a molecule is polar or nonpolar, can have a big effect on the behavior of the molecule.


Nonpolar Covalent BondsNonpolar Covalent Bonds

• Occurs between two atoms of the same electronegativity.

• Electrons are shared equally – Both atoms are pulling with the same force

• Examples – eneg = electronegativity– O=O (O2)

• Same atom – same eneg

– C—H • Carbon and hydrogen have the same eneg


Polar Covalent BondsPolar Covalent Bonds

• Occurs between two atoms of differing eneg• Electrons are not shared equally

– i.e. e- spend more time around one atom than the other

• This creates a slight polarity of charge in the molecule– More eneg atom gains slightly negative charge– Less eneg atom gains a slightly positive charge

• Note – oxygen is the big one here• Example

– H – O bond– Hydrogen is slightly positive, oxygen slightly negative

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Water Molecule (polar)


Electronegativity for Important Electronegativity for Important AtomsAtoms

• F – most eneg

• O – highly eneg

• N – eneg

• Cl – eneg

• C and H are middle of the road eneg

• C – H bond is nonpolar


Ionic BondsIonic Bonds

• Also strong– Relatively weak around water

• Around water, ionically bonded substances dissociate into ions

• Generally occur between a metal and a nonmetal– Metal loses electron, nonmetal gains electron

• Electrons are not shared, they are transferred from one atom to another

• Differences in eneg are great• Ions (charged particles) are formed• An ionic bond is an attraction between oppositely

charged ions.


Examples of Ionic BondsExamples of Ionic Bonds

• Na + Cl Na+ + Cl- NaCl

– Cl steals an e- from Na, gains a 1- charge and leaves Na with a 1+ charge. The oppositely charged ions are attracted.

• Mg + 2F Mg2+ + 2F- MgF2

– Two fluorines steal 1 e- each from Mg, gain a 1- charge and leave Mg with a 2+ charge. The oppositely charged ions are attracted.


Trends for Ionic BondingTrends for Ionic Bonding

• Group IA, 1+ ions– Except H

• Group IIA, 2+ ions• Group VIIA, 1- ions• Group VIA, 2- ions


PracticePractice

• Using LEDD’s . . .– React potassium with iodine– React calcium with chlorine

• Answers– KI

– CaCl2


Hydrogen BondsHydrogen Bonds

• Hbonding is an attraction between the slightly positively charged atom in one polar bond and the slightly negatively charged atom in a different polar bond

• Occur only between polar molecules or polar regions of molecules

• Weak, short-lived bonds (still very important)• This can happen between two different

molecules or between different regions of the same molecule

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Example of a Molecule With Polar Regions (phospholipid)

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Hydrogen Bonding Between Two Water Molecules

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Hydrogen Bonding Between Regions of the Same Molecule

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Hydrogen Bonding in DNA


Van der Waals InteractionsVan der Waals Interactions

• Due to the random movement of electrons• Weak• Short-lived• Can occur in both polar and nonpolar molecules• Only occur when molecules are very close

together• Allows all molecules to be attracted to one

another• Plays role in the shape of larger molecules

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Van der Waals


Molecular Molecular ShapeShape


Molecular ShapeMolecular Shape

• Every covalently bonded molecule has a characteristic size and shape.

FOR IB BIO…the only thing about shape to remember is:

• Biological Structure is related to function– i.e. A molecule’s structure is directly related to

its “job”

• Molecules communicate via shape

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Ethane (C2H6)

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Neurotransmitter Communication

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Cell Surface Receptors

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Enzymes (catalyze reactions)

http://ntri.tamuk.edu/cell/an-enzyme.gif

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Taste

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Saccharine (Sweet ‘n Low)

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Aspartame (Nutra-Sweet)


Chemical ReactionsChemical Reactions


6CO6CO22 + 6H + 6H22O O CC66HH1212OO66 + 60 + 6022

• Represented by chemical equations– Reactants on the left– Products on the right– Some bonds are broken and reformed– Mass is conserved in a reaction

• In a balanced chemical equation, the total # of atoms of each element must be equal on both sides of the equation

• Is this equation balanced?


EquilibriumEquilibrium

• In some reactions, all of the reactants are converted to products

• Most reactions, however, are reversible – they can go in either direction

• CO2 + H2O H2CO3

• Eventually, equilibrium will be met.– This is when the reaction is occurring in both

directions at the same rate


Activation EnergyActivation Energy

• The energy necessary to start a reaction

• Can be high

• This is good – control

• Enzymes (usually proteins) act as catalysts to lower the EA – control


Exergonic/Endergonic Reactions Exergonic/Endergonic Reactions and Free Energyand Free Energy

• Free energy – energy that can be used to do work• Exergonic reactions

– release free energy– result in products with less stored energy than the reactants– Reactants (high E) products (lower E) + E (free)– C6H12O6 + 602 6CO2 + 6H2O + E– Molecules are being broken down (catabolism)

• Endergonic reactions – store free energy– result in products with more stored energy then the reactants– Reactants (lower E) + E (free) products (high E)– 6CO2 + 6H2O + E C6H12O6 + 602

– Molecules are being built up (anabolism)


Oxidation – Reduction ReactionsOxidation – Reduction ReactionsREDOX REDOX

• LEO the lion goes GER

• Loses e- oxidation, gains e- reduction

• Any time an ion is formed – redox reaction

• Example– Na + Cl Na+ + Cl-

• Na has lost e- and has been oxidized

• Cl has gained an e- and has been reduced


Redox in Covalent bondsRedox in Covalent bonds

• Redox rxns can also involve covalent bonding• Atom can be reduced if it becomes bonded to a

highly eneg atom.– i.e. it’s own e- are being pulled away from it

• Example – C-H bond broken, H replaced with O, C-O– Oxygen is highly e-neg– Carbon has been oxidized– Oxygen has been reduced


Dalton’s Atomic TheoryDalton’s Atomic Theory

• We already have discussed this, but to make it more clear the following 5 ideas are the keys to make the Atomic Theory more easy to identify

• 1. Elements are made of tiny particles called atoms.

• 2. All atoms of a given element are identical



• 3. The atoms of a given element are different from those of any other element.

• 4. Atoms of one element can combine with atoms of other elements to form compounds. A given compound always has the same relative numbers and types of atoms.



• 5. Atoms are indivisible in chemical processes. That is, atoms are not created or destroyed in chemical reactions. A chemical reaction simply changes the way the atoms are grouped together.


SOLUTIONSSOLUTIONS


Describing SolutionsDescribing Solutions

• A solution is a uniform mixture• Two types of parts

– Solvent –the dissolving agent• Water is a great example (especially in cells)

– Solutes – are dissolved in the solvent• Anything dissolved in a substance• There can be many solutes in a given solvent

• Example – mix salt and water– Water is the _____– Salt is the _______


Like Dissolves LikeLike Dissolves Like

• Polar vs. nonpolar• Polar and nonpolar substances repel one

another• So . . .• Polar (and ionic) solutes will dissolve in

polar solvents• Nonpolar solutes will dissolve in nonpolar

solvents• Think oil (nonpolar) and vinegar (polar)


Hydrophobic vs. HydrophilicHydrophobic vs. Hydrophilic

• HYDROPHOBIC• Hydro = water• Phobic = fearing• Don’t dissolve in water• Nonpolar substances• OIL

• HYDROPHILIC• Hydro = water• Philic = loving• Do dissolve in water• Polar/ionic substances• VINEGAR


ReviewReview

• Like dissolves like• Hydrophilic and hydrophobic, i.e. nonpolar and

polar molecules, literally repel one another• All polar molecules are hydrophilic• All ionic molecules are hydrophilic• All nonpolar molecules are hydrophobic• However

– Some molecules can be both hydrophobic and hydrophilic (in different areas)

– Example – phospholipids


PHOSPHOLIPIDPHOSPHOLIPIDhydrophobic and hydrophilic regionshydrophobic and hydrophilic regions


Cell MembraneCell Membranephospholipid bilayerphospholipid bilayer


Concentration of a SolutionConcentration of a Solution

• A measure of the amount of solute/solvent

• Lots of solute and/or low solvent = a high concentration (represented by [x] )

• Aqueous solution – water is the solvent– Very important to life

• Saturated solution – cannot dissolve any more solute


Ionic Substance DissolvingIonic Substance Dissolving


Covalent Substance DissolvingCovalent Substance Dissolving


Acids and BasesAcids and Bases


Dissociation into Ions Dissociation into Ions

• To break into separate ions in solution

• Ionically bonded substances do this– NaCl Na+(aq) + Cl-(aq)

• Covalently bonded substances don’t dissociate into ions, with one exception

• Water is the “exception”– H2O H+ + OH-

• Note H+ and H3O+ are synonymous


Acids and BasesAcids and Bases

• Acids

• H3O+ ↔ H+ + H2O

• H3O+ = Hydronium

• Acidity or alkalinity (bases) is actually a measure of hydronium and hydroxide ions dissolved in a solution

• BASES• OH- = hydroxide ion• REMEMBER: NaOH ↔ Na+ + OH-


AcidsAcids

• Proton donors

• Increase H+ (proton) concentration

• Can be strong or weak

• Example of a strong acid– Hydrochloric acid (HCl)– HCl H+ + Cl-

– Dissociates completely


Bases (alkaline)Bases (alkaline)

• Proton acceptors• Decrease H+ (proton) concentration• Can be strong or weak• Example of a strong base

– Sodium hydroxide (NaOH)– NaOH Na+ + OH-

– Dissociates completely– Makes lots of hydroxides which “eat up” protons

– OH- + H+ H2O


pH Scale pH Scale


pH and lifepH and life

• Control of pH is very important to living things (homeostasis)

• Example– Human blood pH range generally 7.35 – 7.45– Anything below 7 or above 7.8 can be deadly

• Buffers– Weak acid/base that can neutralize small amounts of

another acid/base– H2CO3 H+ + HCO3

-

– Carbonic acid hydrogen ion + bicarbonate ion

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