the development of atomic models democritus was a pre- socratic greek philosopher (born around 460...

The Development of Atomic Models

• Democritus was a pre-Socratic Greek philosopher (born around 460 BC).

• Democritus was originator of the belief that all matter is made up of various imperishable, indivisible elements which he called "atomos", from which we get the English word atom.

•According to legend, Democritus was supposed to be mad because he laughed at everything, and so he was sent to Hippocrates to be cured. Hippocrates pointed out that he was not mad, but, instead, had a happy disposition. That is why Democritus is sometimes called the laughing philosopher.

BB - ModelDalton’s ModelMore to come

Plum Pudding Model

• Proposed by J. J. Thomson (1856 - 1940), the discoverer of the electron in 1897.

• The plum pudding model was proposed in March, 1904 before the discovery of the atomic nucleus.

• In this model, the atom is composed of electrons surrounded by a soup of positive charge to balance the electron's negative charge, like plums surrounded by pudding. The electrons were thought to be positioned throughout the atom.

• Electrons could move like letters in alphabet soup

• Instead of a soup, the atom was also sometimes said to have had a cloud of positive charge.

Thomson's model was compared (though not by Thomson) to a British treat called plum pudding, hence the name. It has also been called the chocolate chip cookie model, but only by those who have not read Thomson's original paper

Nuclear Model• The Gold foil experiment or

the Rutherford experiment was an experiment done by Ernest Rutherford (1871 - 1937) in 1909. This experiment discovered the nucleus.

• Led to the downfall of the plum pudding model of the atom.

• Alpha particles (positive particles--Helium Nuclei) were shot at gold foil.

• Particles passed through the gold foil. A few shot back.

• Conclusions:

1. Atom is mostly empty space2. Dense center called the nucleus3. Electrons were stuck surrounding the nucleus.

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Planetary Model• Introduced by Niels Bohr,

a Danish physicist (1885 - 1962), in 1913.

• Because of its simplicity, the Bohr model is still commonly taught to introduce students to quantum mechanics.

• The Bohr model depicts the atom as a small, positively charged nucleus surrounded by waves of electrons in orbit — similar in structure to the solar system, but with electrostatic forces providing attraction, rather than gravity.

"The opposite of a correct statement is a false statement. But the opposite of a profound truth may well be another profound truth." Niels Bohr

Quantum Mechanical Model

• Erwin Schrödinger (August 12, 1887 – January 4, 1961)

• An Austrian physicist, achieved fame for his contributions to quantum mechanics, especially the Schrödinger equation, for which he received the Nobel Prize in 1933.

• This model is based on probability

• Where are you going to find and electron 90% of the time.

• Atom is viewed as a fuzzy cloud.

• Schrödinger equations create electron clouds (orbitals) with specific shapes.

Main Points For “Atoms” Video

• What is the key to understanding atomic structure?• The discovery of what particle is associated with the

Crook’s Tube?• What did Rutherford expect to happen in the gold foil

experiment?• What was Rutherford’s genius?• What conclusions did Rutherford draw from the Gold

foil experiment?• How much smaller is the nucleus than the electron

cloud?• What determines the shape of the electron cloud?

Dalton’s Atomic Theory

• Democritus - Greek philosopher who first suggested atoms

• John Dalton (1766-1844)

• Studied ratios in which elements combine

• Dalton put together the first atomic theory

Dalton’s Atomic Theory• All elements are composed of tiny indivisible particles called

atoms• Atoms of the same element are identical. The atoms of any one

element are different from those of any other element• Atoms of different elements can physically mix together or can

chemically combine with one another in simple-whole number ratios to form compounds

• Chemical reactions occur when atoms are separated, joined, or rearranged. Atoms of one element, however, are never changed into atoms of another element as a result of a chemical reaction

Finding how many subatomic particles for each atom

• Atomic Number - whole number on p.t.

• Gives the number of protons

• Atoms are electrically neutral; there, positives equal negatives.

• Atomic number also equals number of electrons

Finding how many subatomic particles for each atom

• Mass number = protons plus neutrons– Mass number = p+ + no

• Mass number is not found on the periodic table

• So, nO = mass number - p+

• If carbon has a mass number of 14, how many e-’s, p+’s, and no’s does it have?

Symbols

• Mass number is in top left and atomic number is in bottom left

• 16O 9Be 8 4

How many subatomic particles in each?

Oxygen -

Beryllium -

Isotopes

• Atoms with the same number of protons but different numbers of neutrons

• Ex) Carbon 12 vs. Carbon 14

• These atoms have a different mass

• Chemically alike because still have the same number of protons

Isotopes of Hydrogen

• Hydrogen -1 simply called hydrogen

• Hydrogen - 2 called deuterium

• Hydrogen - 3 called tritium

Development of AMUs

• Atomic Mass Units (AMUs)

• Protons have a mass of 1 amu

• 1.67 x 10-24 g

• Neutrons have a mass of 1 amu

• Electrons have a mass of 0 amu

• 9.11 x 10-28 g

Atomic Mass

• The weighted average mass of the isotopes in a naturally occurring sample of the element

• Don’t confuse with “mass number”• To calculate atomic mass you need 3 pieces

of information• 1. The number of stable isotopes• 2.The mass of each isotope• 3.The natural percent abundance of each

isotope

Atomic Mass

• Example Problem - Calculate the atomic mass for element X. One isotope has a mass of 10 amus (10X) and is 20% abundant. The other has a mass number of 11 amus (11X) and an abundance of 80%.

• To solve: Multiply the mass number times the abundance than add them together.

Atomic Mass

• 10 x 0.20 = 2.0

• 11 x 0.80 = 8.8

• Add 2.0 + 8.8 = 10.8– The atomic mass of element X is 10.8

amus

Atomic Mass

• Your turn. Solve:– What is the atomic mass of Element Z?

The isotopes are 16Z, 17Z, 18Z; with percent abundances of 99.759, 0.037, 0.204.

Atomic Mass

• Answer

Atomic (Hotels) Theory• Floors of the hotel are known as energy levels• The period corresponds to the floor. The number of

the floor is called the Principle Quantum Number• Energy levels are divided into sublevels. These are the

rooms of our atomic hotel. There are different types of rooms.

• Atomic Orbital - Region of high probability of finding an electron

• There are 4 types of rooms (orbitals). s, p, d, f rooms• s - Spherical p - dumbbell d- clover leaf f - too

complex to describe. • s - superior p - preferred d- desirable f- fair

– Sharp principal diffuse fundamental

Atomic (Hotels) Theory

• Three Managers each with their own rule:• 1. Aufbau principle • Rooms closest to the basement are checked

out first• Electrons enter orbitals of lowest energy first• 2. Pauli Exclusion Principle - no more than 2

per room• An atomic orbital contains a maximum of 2

electrons

Atomic (Hotels) Theory

• 3. Hund’s Rule• Best rooms are check out first• Similar rooms must have an occupant before pairing

up as roommates• Roommates must have opposite spins to share a

room• When electrons occupy orbitals of equal energy, one

electron enters each orbital until all the orbitals contain one electron with parallel spins

S-orbitals of the 1st and 2nd energy levels.

P-orbital

2nd energy level

The anatomy of the periodic table

• Get out your periodic tables• Know where the following are on your periodic table (p.t)• Group or Family• Period• Group A (representative elements)• Group B• Metals• Nonmetals• Metalloids (Semimetals)

– Note - aluminum is not considered a metalloid

The anatomy of the periodic table

• Know where the following are on your periodic table (p.t) continued

• Transition metals• Inner transition metals• Alkali metals• Alkaline metals• Halogens• Noble gases • Atomic Number• Atomic Mass

Focus Questions to upcoming video8 Questions

• What is the periodic law?• Who is Dmitri Mendeleev and what did he do?• How is today’s periodic table different from Mendeleev’s?• What characteristic is common among the noble gases• What are the names for vertical columns and horizontal

rows• What characteristics are common amongst group 1A?• What are the 2 most important things about the periodic

table?• Why is fluorine the Tyrannosaurus rex of the periodic

table?

Atoms

10-15 m

Neutron

Proton

Nucleus (protons and neutrons)

Space occupied by electrons

10-10 m

Periodic Table and Electron Configurations

• Build-up order given by position on periodic table; row by row.

• Elements in same column will have the same outer shell electron configuration.

The relation between orbital filling and the periodic table

Electron Configuration

• Orbitals have definite shapes and orientations in space

(insert Fig 2.11 of text)(if it will not all fit on one screen, put part (a) on one screen and part (b) on the next)

Orbital occupancy for the first 10 elements, H through Ne.

Atomic radii of the main-group and transition

elements.

Trends in the Periodic Table

Trend for atomic radii

• Left to right atoms get smaller• Why?

– Increase in nuclear charge– More protons and more electrons means greater

electrostatic attractions (stronger magnet)

• Top to bottom atoms get larger• Why?

– Increase in energy levels (You are adding floors to your hotel). Electrons are further from the nucleus

Atomic Radius• Atomic radii actually

decrease across a row in the periodic table. Due to an increase in the effective nuclear charge.

• Within each group (vertical column), the atomic radius tends to increase with the period number.

Atomic Radii for Main Group Elements

Trend for Ion Size• Ion is a charged atom. • Metals lose electrons and nonmetals gain electrons to create

ions. • Cations are pawsitive (positive) and Anions are negative.• Cations are smaller than their corresponding atom. Why?• Loss of electrons means the positive nucleus has a greater

attraction on the remaining electrons• Anions are larger than their corresponding atom. Why?• Gain of electrons means the nucleus has less attraction for the

electrons as well as the electrons are repulsing each other causing an increase in the size of the electron clouds

This is a “self-consistent” scale based on O-2 = 1.40 (or 1.38) Å.

Ionic radii depend on the magnitude of the charge of the ion and its environment. (more later)

Positively charged ions are smaller than their neutral analogues because of increased Z*.

Negatively charged ions are larger than their neutral analogues because of decreased Z*.

Radii of ions

Same periodic trends as atomic radii for a given charge

Trend for ion size

• Decrease across a period then jumps in size at nonmetals and continues to decrease

• Increases on the way down a group as you are adding energy levels (electrons are farther from the nucleus)

Ionization energy

• The energy required to remove an electron

First ionization energies of the

main-group elements

Trends in the Periodic Table

Ionization Energy• Ionization energy is a periodic property

• In general, it increases across a row. Why?• increasing attraction as the number of protons in the

nucleus increases (stronger magnet)• it decreases going down a group. Why?• Outer shell electrons are further from the nucleus so

less electrostatic attraction. Nucleus has less pull on them.

• Shielding also plays a factor.

Ionization energy

6) The trend across from left to right is accounted for by a) the increasing nuclear charge.

Electronegativity - This is the most important trend to understand for this class.

• The tendency for an atom to attract electrons when chemically bonded.

• Same trend as ionization energy.– In general, it increases across a row. Why?– increasing attraction as the number of protons in

the nucleus increases (stronger magnet)– it decreases going down a group. Why?– Outer shell electrons are further from the nucleus

so less electrostatic attraction. Nucleus has less pull on them. Shielding also plays a factor.

Trends in three atomic properties

See chart in book for summary

Check for understanding

• Which of the following atoms has the largest atomic radii, ion size, electronegativity, and ionization energy

• Na, Mg, K, Ca, S, Cl, Se, Br

Bonding

Ionic

Metallic

Covalent

Valence Electrons

• Atoms in a group behave similarly because they have the same number of valence electrons.

• Valence electrons - electrons in the highest occupied energy level

• To find the number of valence electrons just look at the group number

Lewis Electron Dot Structures

• Symbol of element with dots around it representing valence electrons

• Example:

C 2 electrons per side totaling 8

No pairs until each room has an electron

Lewis Electron Dot Structures

• Example: O• Pairs of electrons are adjacent not

across from one another. This will help with identifying shape.

Unshared Pair or Lone Pair

Octet Rule

• In forming compounds, atoms tend to achieve the noble gas electron configuration.

• They lose, gain, or share electrons with another atom to achieve 8.

• Metals lose electrons leaving a complete octet in the lower energy level

• Nonmetals gain electrons to fill the energy level to achieve 8.

Electron Configurations of Ions

• Na - 1s22s22p63s1 Na1+ - 1s22s22p6

• O - 1s22s22p4 O2- - 1s22s22p6

• Write the following on your periodic table.• Group 1A - 1+ - loses 1 electron• Group 2A - 2+ - loses 2 electrons• Group 3A - 3+ - loses 3 electrons• Group 4A - - Depends on the atom• Group 5A - 3- - gains 3 electrons• Group 6A - 2- - gains 2 electrons• Group 7A - 1- - gains 1 electron• Group 8A (0) - does not form ions

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VSEPR Theory

• Valence Shell Electron Pair Repulsion Theory

• “Electron pairs around atoms tend to be as far apart as possible.”

• Similar charges (I.e., negative charges from electrons) tend to repel each other and want to be spaced apart at maximum angles.

• Used to predict molecular geometries

• Bond angles

– Angles between bonds

– Spacing apart as far as possible

• Lone pairs of electrons will repel bonded atoms a bit more than expected toward each other around the central atom

Covalent Bonding

Polar Bonds and Molecules

Covalent Bonding-- Polar Bonds and Molecules --

Bond Polarity• “The Tug of War”

– The pairs of electrons that are bonds between atoms are pulled between the nuclei of the atoms in a bond.

– The electronegativities of the atoms determines who is winningYet; there is no winner. The tug of war never ends.

• Classifications for Bonds– Nonpolar covalent

• When atoms pull the bond equally• Happens with two atoms of equal electronegativity, most often

using the same atoms• Examples: H2, O2, N2

– Polar covalent• When atoms pull the bond unequally• Happens with two atoms of different electronegativities• Example: HCl, HF, NH


Bond Polarity

• In a polar molecule, one end of the molecule is slightly more electronegative than the other atom, resulting in one atom being slightly negative (-) because of higher electronegativitiy, and the other atom being slightly positive (+) because of lower electronegativity.

is known as a partial charge since it is much less than 1+ or 1- charge.


Bond Polarity• Electronegativities and Bond Types

– H: 2.1 Cl: 3.0 Since hydrogen is less, it will have the positive partial charge while chlorine has the negative partial charge.

– 3.0 – 2.1 = 0.9 HCl is polar covalent.

0.0 – 0.1 difference Nonpolar covalent bond H – H (0.0 difference)

0.1 – 1.7 difference Polar covalent bond H – Cl (0.9 difference)

1.7 + difference Ionic bond Na+Cl- (2.1 difference)


Polar Molecules• Dipole

– Molecule that has two poles– Example: HCl from the previous page

• Polar vs. Nonpolar

H2O and CO2

Both have 3 atoms; yet,

One is polar and one is nonpolar.

Why?

Structure (with bond polarity) determines the

molecules polarity.

3 video clips coming up

• Get out your bonding sheet of chemistry• Yes, we are going to discuss more on bonding so try

and hold your enthusiasm as rioting is not tolerated and things need to be accomplished.

• Bonding appreciates your cooperation and will sign autographs at the end of the period.

• Thank you.

Intermolecular Attractions

Attractions Between Molecules• van der Waals forces

– Two types: dispersion forces and dipole interactions• Dispersion forces

– Weakest of all molecular interactions– Caused by movement of electrons– Occurs in the BrINClHOF’s

Intermolecular Attractions Attractions Between Molecules

2nd van der Waals force• Dipole interactions

• Occurs when polar molecules are attracted to one another

• Partial charge (+) of one polar molecule is attracted to the opposite partial charge (-) of another molecule

Intermolecular Attractionsattractions between molecules

• Hydrogen bonding– Hydrogen covalently bonded to a very electronegative atom is also

weakly bonded to an unshared electron pair of another electronegative atom

– Example: water

Short Lab - No lab report

• In the back in a tray is a micropipet and a penny. Place 1 drop of water on the penny and then take a guess as to how many drops of water you can fit on a penny. Write your guess on the board at the front of the room. Place your name next to your guess. Then count the drops of water you fit on the penny before it overflows. Record this count on the board next to your guess. As your water drop grows, watch it from the side. Clean up and have a seat at your desk.

Gases

• There are four variables that affect a gas.

• 1. Pressure

• 2. Volume

• 3.Temperature

• 4. Number of molecules

The variables

• Pressure units - there are many units for pressure. – kPa - kilopascal (101.3)– Atm - atmospheres (1 )– mm Hg - millimeters of Hg (760) - torr

• Volume is measured in Liters• Temperature is in Kelvin • K = oC + 273 or oC = K - 273• If the Kelvin temperature doubles the K.E. doubles.

The pressure-volume relationship Boyle’s Law

• Pressure and volume are inversely related.

• One goes up the other goes down

• P1 x V1 = P2 x V2

Boyle’s Lawsample problem

• A high-altitude balloon contains 30.0 L of helium gas at 103 kPa. What is the volume when the balloon rises to an altitude where the pressure is only 25.0 kPa?

• 103 kPa x 30.0 L = 25.0 kPa x V2

• V2 = 124 L


• Your turn

• Your birthday balloon travels with you from Lincoln to Denver. The balloons volume is 4.0 L with an atmospheric pressure of 101.3 kPa. You arrive in Denver where the atmospheric pressure is 90.0 kPa. What is the new volume of your balloon?


• Answer

The temperature-volume relationshipCharles’s Law

• Volume and temperature have a direct relationship.

• One goes up the other goes up– One goes down the other goes down

V1 = V2

T1 T2

Charles’s Lawsample problem

• A balloon inflated in a room at 24oC has a volume of 4.00 L. The balloon is then heated to a temperature of 58oC. What is the new volume if the pressure remains constant?

4.00L = V2

297 K 331 K

V2 = 4.46 L


• Your turn

• A balloon is inflated in a room at 24oC and has a volume of 4.00 L. The balloon is placed in a freezer and then removed the volume is now 3.25 L. What was the temperature of the freezer in Celsius?


• Answer

The Temperature-Pressure RelationshipGay-Lussac’s Law

• The pressure of a gas is directly proportional to the temperature of a gas

• Temperature goes up; pressure goes up

P1 = P2

T1 T2

Gay-Lussac’s Lawexample problems

• The gas left in a used aerosol can is at a pressure of 103 kPa at 25oC. If the can is thrown into a fire, what is the pressure of the gas when it reaches 928oC?

103 kPa = P2

298 K 1201 K

P2 = 415 kPa


• Your turn

• A container of propane has a pressure of 108.6 kPa at a morning temperature 15oC. By mid afternoon the temperature has reached 32oC. What is the pressure inside the propane tank?


• Answer

The combined gas law

P1 x V1 = P2 x V2

T1 T2

The combined gas lawexample problem

• The volume of a gas-filled balloon is 30.0 L at 40oC and 153 kPa. What volume will the balloon have at STP?

153 kPa x 30.0 L = 101.3 kPa x V2

313 K 273 K

V2 = 39.5 L


• Your turn

• A gas-filled balloon is 25.0 L at 35oC and 145 kPa. What is the temperature if the volume increases to 28.0 L and a pressure of 152 kPa?


• Answer

The Ideal Gas Law

PV=nRT• P = Pressure

• V = Volume

• n = Number of Moles

• R = ideal gas constant

• T = temperature in Kelvin

R -The ideal gas constant

• Depends on unit of pressure

• 0.0821 L . Atm / K . mol

• 62.4 L . mmHg / K . mol (torr is mm Hg)

• 8.31 L . kPa / K . mol

Ideal Gas Law example problem

• Calculate the pressure of 1.65 g of helium gas at 16.0oC and occupying a volume of 3.25 L?

• You will need g to moles and Celsius to Kelvin:• 1.65 g He 1 mole He• 4.0 g He = 0.413 mol He• K = oC + 273 ; 16. 0 + 273 = 289 K• For this problem you will need to pick an R value. For this

problem I will choose to use the R value containing kPa. I picked it. You can’t do anything about it. So; just try and stop me.

• Plug and Chug baby, get ‘R done. Do it. Come on I dare ya.• Get it - ‘R as in ideal gas constant


• P x 3.25 L = 0.413 mol x 8.31 kPa . L x 289 K• mol . K• Do the algebra and solve; if you do it right, guess

what? You get the answer right. Neat concept, huh? Maybe your mommy will give you a cookie.

• = 305 kPa• Your turn• How many moles of gas are present in a sample of

Argon at 58oC with a volume of 275 mL and a pressure of 0.987 atm.


• Answer• 0.987 atm x 0.275 L = n x 0.0821 L . Atm x 331K• mol . K• Do the dew; oops, I mean the algebra and presto; the answer

with the correct number of sig figs is..• Do you know how to keep a so called chem student in

suspense?• Do ya?• Do ya?• = 0.00999 mol Ar• Congrats - you can plug and chug.• Bye Bye now.

Surface Tension

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Intermolecular Forces Bulk and Surface

Phase ChangesPhase ChangesEnergy Changes Accompanying Phase ChangesEnergy Changes Accompanying Phase Changes• All phase changes are possible under the right

conditions (e.g. water sublimes when snow disappears without forming puddles).

• The sequence

heat solid melt heat liquid boil heat gas

is endothermic.• The sequence • cool gas condense cool liquid freeze cool solid

is exothermic.

Phase ChangesPhase Changes

Heating CurvesHeating Curves

Heating Curve Illustrated

Vapor PressureVapor Pressure

Explaining Vapor Pressure on Explaining Vapor Pressure on the Molecular Levelthe Molecular Level

• Dynamic Equilibrium: the point when as many molecules escape the surface as strike the surface.

• Vapor pressure is the pressure exerted when the liquid and vapor are in dynamic equilibrium.

Vapor PressureVapor Pressure

Volatility, Vapor Pressure, and Volatility, Vapor Pressure, and TemperatureTemperature

• If equilibrium is never established then the liquid evaporates.

• Volatile substances evaporate rapidly.

• The higher the temperature, the higher the average kinetic energy, the faster the liquid evaporates.

Liquid Evaporates when no Equilibrium is Established

Vapor PressureVapor PressureVapor Pressure and Boiling PointVapor Pressure and Boiling Point• Liquids boil when the external pressure equals the vapor

pressure.• Temperature of boiling point increases as pressure

increases.• Two ways to get a liquid to boil: increase temperature or

decrease pressure.– Pressure cookers operate at high pressure. At high

pressure the boiling point of water is higher than at 1 atm. Therefore, there is a higher temperature at which the food is cooked, reducing the cooking time required.

• Normal boiling point is the boiling point at 760 mmHg (1 atm).

Thermochemistry-- The Flow of Energy: Heat --

Thermochemistry: the study of heat changes in chemical reactions

Chemical potential energy: energy stored within the structural units of chemical substances


Chemical System Types

System type Descriptionq (change

in heat)

EndothermicSystem absorbing heat from the surroundings

q > 0

ExothermicSystem releasing heat to

the surroundingsq < 0


Law of Conservation of Energy:

In any chemical or physical process, energy is neither created nor destroyed


The calorie• Expressed as a c (lower case)• Quantity of heat needed to raise the temperature

of 1 g of pure water 1CCalorie• Expressed as a C (upper case)• Dietary Calorie• 1 Calorie = 1 kilocalorie = 1000 calories

Thermochemistry-- An Intro Video --



Joule• SI unit of heat and energy• Raises the temperature of 1 g of pure water 0.2390C• 4.184 J = 1 calHeat Capacity• Amount of heat needed to increase the temperature of an object

exactly 1C• Will change depending on the mass and chemical compositionSpecific Heat

Quantity of heat needed to raise the temperature of 1g of substance 1oC


Specific Heat Capacity

Heat (q) specific heat capacity (C)

Mass (m) change in temperature (T)

q = mC T


Example:How many kilojoules of heat are absorbed when 1.00 L of water is heated from 18C to 85C?Solution:

q = mCTq = 1000g x 4.18 J x 67oC

g o C

q = 2.8E5 J 1 KJ 1000 J

= 280 KJ


Example:

A chunk of silver has a heat capacity of 42.8 J/C. If the silver has a mass of 181 g, calculate the specific heat of silver.

Solution:

q = mCT

42.8 J = 181g x C x 1OC

C = 0.236 J/goC

Thermochemistry-- Measuring and Expressing Heat Changes --

Your Turn:The temperature of a piece of copper with a mass of

95.4 g increases from 20.0oC to 43.0oC when the metal absorbs 849 J of heat. What is the specific heat of copper?

Thermochemistry-- Measuring and Expressing Heat Changes --

Calorimeter

Properties of acids and Bases• Taste

– Acids taste sour ex - lemons– Bases taste bitter ex - soap

• Feel– Acids feel like water; but have you ever

gotten fruit juice on a canker sore or cut– Bases feel slippery ex - soap and water

Properties of acids and bases• Reaction with metals

– Acids - Hydrogen gas is produced when reacted with certain metals

– Bases - typically don’t react with metals

• Both are electrolytes

• React to form salt and water

• Milk of Magnesia (Magnesium Hydroxide) is a base used to treat excess stomach acid problems.

Definitions of Acids and Bases• Arrhenius Acid/Base - focused on

products– HCl + H2O --> H3O+ + Cl-

– NaOH + H2O --> Na+(aq) + OH-

(aq)

• Acids form H+’s and Bases form OH-’s

Definitions of Acids and Bases• Bronsted-Lowery Acid/Base - focused

on what happens during formation– HCl + H2O --> H3O+

(aq) + Cl-(aq)

• Acid - substance that donates a proton

• Base - substance accepts a proton

• In the above example, what is the base and what is the acid?

• How about this example?– NH3 + H2O --> NH4

+(aq) + OH-

(aq)

Definitions of Acids and Bases• Lewis acid/base

– H+ + OH- ---> H2O

• Acid - a substance that can accept a pair of electrons to form a covalent bond

• Base - a substance that donates a pair of electrons to form a covalent bond

• What is the Lewis acid and base?– AlCl3 + Cl- --> AlCl4-

Problem

• Write an equation for the ionization of nitric acid and explain how it fits each definition?

• HNO3 --> H+ + NO3- or

• HNO3 --> H3O+ + NO3- or

• HNO3 + H2O --> H3O+ + NO3-

Hydrogen Ions from WaterThis is the “bases” to start understanding pH

• Water is considered neutral• Collision between water molecules can cause

a hydrogen ion to transfer from one molecule to another.

H2O H2O H3O+ OH-

hydronium ion hydroxide ion

Self-Ionization of Water

Water self ionizes to the concentration of 1.0 x 10-7 mol/L.

When the concentration of each ion equals 1.0 x 10-7 mol/Lthe solution is said to be neutral

Therefore, since water is considered neutral the concentrations of the ions can be calculated through the Ion product constant.

The Ion-Product ConstantKw

• Notation– [H+] - concentration of hydrogen ions

• Or hydronium ions– [OH-] - concentration of hydroxide ions

• When [H+] and [OH-] are multiplied we get the ion-product constant.

• Kw = [H+] x [OH-] = 1.0 x 10-14 (mol/L)2 or M2

• This is an inverse relationship.– One goes up, the other goes down.

The Ion-Product ConstantKw example problem

• If [H+] = 1.0 x 10-5 mol/L, is the solution acidic, basic, or neutral? What is the [OH-] of this solution?

• Answer– Acidic - the [H+] is greater than 1.0 x 10-7

mol/L– 1.0 x 10-5 mol/L x [OH-] = 1.0 x 10-14 M2

– [OH-] = 1.0 x 10-9 mol/L

The Ion-Product ConstantKw example problem

• If [OH-] = 2.8 x 10-8 mol/L, is the solution acidic, basic, or neutral? What is the [H+] of this solution?

• Answer

The pH concept

• [H+] is cumbersome so the pH scale was created.

• pH is the negative logarithm of the hydrogen-ion concentration.

• pH = -log[H+]

Sample pH problems1 of 3

• The hydrogen-ion concentration of a solution is 2.7 x 10-10 mol/L. What is the pH of the solution?

• Answer


• The pH of a solution is 6.8. What is the [H+]?

• Answer


• What is the pH of a solution if the [OH-] = 4.0 x 10-11 mol/L

• Answer

Naming Acids

And writing formulas

General Form - HX (X is an anion or polyatomic ion)

• Rules - 3 of them

• 1. When anion ends in -ide, acid name begins with hydro and -ide is changed to -ic with the word acid

• Ex - HCl is hydrochloric acid

• Try - H2S

• Rule 2 - • When anion ends in -ite, ending changes to -

ous with the word acid. (No hydro)• Name these - H2SO3 , HNO2

• Rule 3 -• When anion ends in -ate, ending changes to -

ic with the word acid• Name these - HNO3 , HC2H3O2

Work backwards to get the formula

• Write the formula for the following:– Chloric acid– Hydrobromic acid– Phosphorous Acid

• Don’t confuse phosphorous and phosphorus

• Answers

Neutralization Reaction

• Base and acid react to produce a salt and water

Titration

• A lab technique where a neutralization reaction is performed to determine the concentration of an unknown.

The anatomy of a titration:

• Standard solution - solution of known concentration.

• End Point - the point at which neutralization is achieved

• Indicator - chemical that changes color with a change in pH. We will be using phenolphthalein.Clear in an acid pink in a baseYou want a light pink

• Buret - Measurement device

Titration Calculations4 steps

• Start with the balanced equation

• Find the moles in the standard solution

• Set up ratio to find moles of unknown

• Find molarity(mol/L) or volume

Titration Calculations example problem

• A 25.75 ml solution of H2SO4 is neutralized by 18.23 ml of 1.0 M NaOH. What is the concentration of H2SO4?

• H2SO4 + 2 NaOH --> Na2SO4 + 2 H2O

• 0.01823 L NaOH 1.0 mol NaOH 1 mol H2SO4

1 L NaOH 2 mol NaOH 0.02575L H2SO4

• = 0.35 mol/L H2SO4

Titration Problemsyour turn

• What is the molarity of phosphoric acid if 15.0 mL of the solution is neutralized by 8.5 mL of 0.15 M NaOH?

Titration Problems1 more your turn

• How many milliliters of 0.45 M hydrochloric acid must be added to 25.0 mL of 0.15 M NaOH?

the development of atomic models democritus was a pre- socratic greek philosopher (born around 460...

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