Atomic Structure Topic 1 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

Atomic - Molecular Theory of Matter The Atomic - Molecular Theory of Matter states that all matter is composed of small, fast moving particles called atoms. These atoms can join together to form molecules. This theory is really thousands of individual theories that provide evidence for the whole theory. This theory is really thousands of individual theories that provide evidence for the whole theory.

Where did it all begin? The word “atom” comes from the Greek word “atomos” which means indivisible. The word “atom” comes from the Greek word “atomos” which means indivisible. The idea that all matter is made up of atoms was first proposed by the Greek philosopher Democritus in the 5th century B.C. The idea that all matter is made up of atoms was first proposed by the Greek philosopher Democritus in the 5th century B.C.

Atoms The smallest particle of an element that retains the properties of that element.

Dalton’s Atomic Theory (1808) 1. Elements are composed of extremely small particles called atoms. 3. Atoms of different elements can combine in simple whole number ratios to form compounds. 4. Chemical reactions only involve the rearrangement of atoms. Atoms are not created or destroyed in chemical reactions All atoms of a given element are identical. The atoms of one element are different from those of any other element

8 X 2 Y 16 X8 Y + 2.1

2

Electrons, protons & neutrons Which one of Dalton’s assumptions is wrong?

Electron negatively charged subatomic particle Sir Joseph J Thomson experimented with cathode rays He found (1897) that cathode rays could be deflected by electrically charged plates, towards (+) side plate J.J. Thompson J.J. Thompson

J.J. Thomson, measured mass/charge of e - (1906 Nobel Prize in Physics) 2.2

e - charge = x C Thomson’s charge/mass of e - = x 10 8 C/g e - mass = 9.10 x g Measured mass of e - (1923 NobelPrize in Physics) Millikan's Oil Drop Experiment 2.2

(Uranium compound) 2.2

Proton 1840x heavier than an electron Approx 1AMU (atomic mass unit) A hydrogen atom stripped of an electron (+) charged particle

Neutron Discovered in 1932 by Chadwick No charge Mass is equal to that of a proton

Chadwick’s Experiment (1932) H atoms - 1 p; He atoms - 2 p mass He/mass H should = 2 measured mass He/mass H = 4  + 9 Be 1 n + 12 C + energy neutron (n) is neutral (charge = 0) n mass ~ p mass = 1.67 x g 2.2

Subatomic Particles mass p = mass n = 1840 x mass e - 2.2

Structure of the Atom Rutherford discovered the nucleus of the atom The nucleus is composed of protons and neutrons

HISTORY OF THE ATOM 1910 Ernest Rutherford oversaw Geiger and Marsden carrying out his famous experiment. they fired Helium nuclei at a piece of gold foil which was only a few atoms thick. they found that although most of them passed through. About 1 in 10,000 hit 1910

1.atoms positive charge is concentrated in the nucleus 2.proton (p) has opposite (+) charge of electron 3.mass of p is 1840 x mass of e - (1.67 x g) 4.TeacherTube - Rutherford\'s experimentTeacherTube - Rutherford\'s experiment  particle velocity ~ 1.4 x 10 7 m/s (~5% speed of light) (1908 Nobel Prize in Chemistry) 2.2

Nucleus (+ ) charge Occupies a small volume Contains protons and neutrons High density

atomic radius ~ 100 pm = 1 x m nuclear radius ~ 5 x pm = 5 x m Rutherford’s Model of the Atom 2.2

Outside the nucleus Electron are found outside the nucleus Electrons occupy most of the volume

Is this really an Atom? The model above represents the most modern version of the atom. The model above represents the most modern version of the atom. (Artist drawing) (Artist drawing) Many of the models that you have seen may look like the one below. It shows the parts and structure of the atom. Even though we do not know what an atom looks like, scientific models must be based on evidence.

How can Indirect Evidence be Gathered? Click here to visit a lab where actual scientific research on atoms is conducted. Click here to visit a lab where actual scientific research on atoms is conducted.Click here to visit a lab where actual scientific research on atoms is conducted.Click here to visit a lab where actual scientific research on atoms is conducted.

Differences in elements Different number of subatomic particles C 6p +, 6n 0, 6e- F 9p +, 10n 0,9e-

Identifying features of Atoms A. Atomic Number (Z) The number of protons in the nucleus of the atom of a certain element. This identifies the element. The atom is electrically neutral therefore it must also have eight electrons around its nucleus

Period Group Alkali Metal Noble Gas Halogen Alkali Earth Metal 2.4

Atomic Mass Unit (AMU) Mass of the p + and n 0 are very small 1.67 x g This unit is called the Atomic Mass Unit The AMU is 1/12 the mass of a Carbon – 12 atom that contains 6p and 6n

Atomic number (Z) = number of protons in nucleus Mass number (A) = number of protons + number of neutrons = atomic number (Z) + number of neutrons # OF NEUTRONS = mass number – atomic number X A Z H 1 1 H (D) 2 1 H (T) 3 1 U U Mass Number Atomic Number Element Symbol 2.3

ISOTOPES The nuclei of atoms must contain the same # of protons but neutrons may vary

ISOTOPES Atoms that have same # of p+ Different # of neutrons Different mass numbers

2.3

How many protons, neutrons, and electrons are in C 14 6 ? How many protons, neutrons, and electrons are in C 11 6 ? 6 protons, 8 (14 - 6) neutrons, 6 electrons 6 protons, 5 (11 - 6) neutrons, 6 electrons Do You Understand Isotopes? Isotope Maker 2.3

Atomic Mass Calculations Atomic Mass Calculations Atomic Mass Atomic Mass Calculations The weighted average of the masses of the isotopes of that element. Most elements occur as two or more isotopes in nature Similar to your class average

An ion is an atom, or group of atoms, that has a net positive or negative charge. An atom that has lost or gained one or more electrons cation – ion with a positive charge If a neutral atom loses one or more electrons it becomes a cation. anion – ion with a negative charge If a neutral atom gains one or more electrons it becomes an anion. Na 11 protons 11 electrons Na + 11 protons 10 electrons Cl 17 protons 17 electrons Cl - 17 protons 18 electrons 2.5

A monatomic ion contains only one atom A polyatomic ion contains more than one atom 2.5 Na +, Cl -, Ca 2+, O 2-, Al 3+, N 3- OH -, CN -, NH 4 +, NO 3 -

13 protons, 10 (13 – 3) electrons 34 protons, 36 (34 + 2) electrons Do You Understand Ions? 2.5 How many protons and electrons are in Al ? 3+3+ How many protons and electrons are in Se ?

2.5

Some Polyatomic Ions (Table 2.3) 2.7

4S 3S 2S 1S 3P 2P

The Development of Atomic Models The timeline shoes the development of atomic models from 1803 to

Electrons in Atoms Dalton –Indivisible Thomson –Raisin bun model Rutherford –Dense, empty Bohr –Definite ;energy –Circular fixed distance from the nucleus –Greater; electron –Energy levels or shells

The Development of Atomic Models The timeline shows the development of atomic models from 1913 to

Quanta Photons ; bundles of energy The energy that is absorbed as electrons jump to higher energy levels and it is emitted when they fall to their lower energy levels If electrons remain in their orbit they don’t lose energy

Niels Bohr and The Planetary Model of the Atom Niels Bohr and The Planetary Model of the Atom All atoms that are in the lowest energy level are in their normal state or ground state b. Principle Energy levels Electrons normally occupy the lowest energy levels. If an atom absorbs energy from an outside source, it may cause the electrons to move to a higher energy levels. This is called the excited state This state is unstable.

Principle Energy level The PEL denotes how far the electron is from the nucleus K- 1____________ L-2____________ M-3____________ Similar to rungs on a ladder INCREASING PE

The Bohr Model Like the rungs of the strange ladder, the energy levels in an atom are not equally spaced. The higher the energy level occupied by an electron, the less energy it takes to move from that energy level to the next higher energy level. 5.1

Spectral Lines When electrons in an atom are in the excited state and return to lower energy levels the energy is emitted as Radiant energy of a specific frequency which produces a characteristic spectral line which can be used to identify elements

Spectra of several elements

+P nm 700 nm 400 nm -e Balmer series for Hydrogen Atom 1 -e -e 5 -e 486 nm 434 nm -e 410 nm

n=1 n=2 n=3 n=4 Spectrum UV IR VisibleVisible Ground State Excited State Excited State unstable and drops back down Energy released as a photon Frequency proportional to energy drop Excited State But only as far as n = 2 this time

BOHR VS ORBITAL MODEL OF THE ATOM It does not represent electrons as moving in planetary orbits around the nucleus Electrons occupy regions of space around the nucleus (not circular paths) Electrons occupy orbitals that may differ in size, shape, or orientation in space

Electron Configurations Ground State The ground state is the most stable energy state of an atom. It is the nature of things to seek the lowest possible energy levels. Therefore, high energy systems are unstable

Electron Configurations The way in which electrons are arranged around the nuclei of atoms 1s 2 2s 2 2p 3 ground state 1s 2 2s 1 2p 4 excited state

Energy levels The energy levels are represented by quantum numbers N is equal to the number of the principle energy level as referred to under the Bohr atom and is the same as the period number in the periodic table (horizontal)

Period Group Alkali Metal Noble Gas Halogen Alkali Earth Metal 2.4

Sublevels The energy levels may be divided into sublevels. Every PEL has one or more sublevels within it. The number of sublevels in the PEL is the same as the principle quantum # PEL 1 has 1 Sublevel s PEL 1 has 1 Sublevel s PEL 2 has 2 Sublevelssp PEL 2 has 2 Sublevelssp PEL 3 has 3 Sublevelss p d PEL 3 has 3 Sublevelss p d PEL 4 has 4 Sublevelsspdf PEL 4 has 4 Sublevelsspdf

Orbitals Each sublevel may consist of one or more orbitals. Only two electrons to occupy each orbital These two electrons have opposite spins

Atomic Orbitals The numbers and kinds of atomic orbitals depend on the energy sublevel. 5.1

Atomic Orbitals Different atomic orbitals are denoted by letters. The s orbitals are spherical, and p orbitals are dumbbell-shaped. 5.1

Atomic Orbitals Four of the five d orbitals have the same shape but different orientations in space. 5.1

Atomic Orbitals The number of electrons allowed in each of the first four energy levels are shown here. 5.1

Electron Configurations –Aufbau Principle According to the aufbau principle, electrons occupy the orbitals of lowest energy first. In the aufbau diagram below, each box represents an atomic orbital. 5.2

3 Rules Pauli Exclusion Principle –An orbital may contain only 2 electrons at the most with opposite spins Hunds Rule –Electrons occupy orbitals of equal energy 1 electron enters each orbital until all orbitals contain one electron with parallel spins

Electron Configurations Orbital Filling Diagram 5.2

What do electron configurations tell us? Principle energy level (PEL) Type of sublevel The number of electrons in the sublevel Atomic Electron Configurations Atomic Electron Configurations 1s 2 PEL Type of sublevel and # of Orbitals # of electrons in orbital

1s 2s 3s 4s 2p 3p 3d Energy Ar 1s 2 2s 2 2p 6 3s 2 3p 6 4p

1s 2s 3s 4s 2p 3p 3d Energy Sc 1s 2 2s 2 2p 6 3s 2 3p 6 3d 1 4s 2 4p

Valence Shell The outer most occupied PEL, the electrons in this shell are called valence electrons an atom can’t have more than 8 electrons in its valence shell ** The chemical properties of an element are determined mainly by the arrangement of electrons in the valence shell**

Kernel The part of the atom including the nucleus that is stripped of its valence electrons

Ground Vs. Excited State Na Na 2 – 7 -2 What element is this? Is it in the ground or excited state?

Ionization Energy (IE) The amount of energy needed to remove the most loosely held electron from a neutral atom Highest – noble gases Metallic – smallest IE Non – Metallic – largest IE

ionic compounds consist of a cation and an anion the formula is always the same as the empirical formula the sum of the charges on the cation and anion in each formula unit must equal zero The ionic compound NaCl 2.6

A molecule is an aggregate of two or more atoms in a definite arrangement held together by chemical bonds H2H2 H2OH2ONH 3 CH 4 A diatomic molecule contains only two atoms H 2, N 2, O 2, Br 2, HCl, CO A polyatomic molecule contains more than two atoms O 3, H 2 O, NH 3, CH 4 2.5

2.6

A molecular formula shows the exact number of atoms of each element in the smallest unit of a substance An empirical formula shows the simplest whole-number ratio of the atoms in a substance H2OH2O H2OH2O molecularempirical C 6 H 12 O 6 CH 2 O O3O3 O N2H4N2H4 NH 2 2.6

Formula of Ionic Compounds Al 2 O x +3 = +63 x -2 = -6 Al 3+ O 2- CaBr 2 1 x +2 = +22 x -1 = -2 Ca 2+ Br - Na 2 CO 3 1 x +2 = +21 x -2 = -2 Na + CO 3 2-

Chemical Nomenclature Ionic Compounds –often a metal + nonmetal –anion (nonmetal), add “ide” to element name BaCl 2 barium chloride K2OK2O potassium oxide Mg(OH) 2 magnesium hydroxide KNO 3 potassium nitrate 2.7

Transition metal ionic compounds –indicate charge on metal with Roman numerals FeCl 2 2 Cl - -2 so Fe is +2 iron(II) chloride FeCl 3 3 Cl - -3 so Fe is +3 iron(III) chloride Cr 2 S 3 3 S so Cr is +3 (6/2)chromium(III) sulfide 2.7

Molecular compounds –nonmetals or nonmetals + metalloids –common names H 2 O, NH 3, CH 4, C 60 –element further left in periodic table is 1 st –element closest to bottom of group is 1 st –if more than one compound can be formed from the same elements, use prefixes to indicate number of each kind of atom –last element ends in ide 2.7

HIhydrogen iodide NF 3 nitrogen trifluoride SO 2 sulfur dioxide N 2 Cl 4 dinitrogen tetrachloride NO 2 nitrogen dioxide N2ON2Odinitrogen monoxide Molecular Compounds 2.7 TOXIC ! Laughing Gas

An acid can be defined as a substance that yields hydrogen ions (H + ) when dissolved in water. HCl Pure substance, hydrogen chloride Dissolved in water (H + Cl - ), hydrochloric acid An oxoacid is an acid that contains hydrogen, oxygen, and another element. HNO 3 nitric acid H 2 CO 3 carbonic acid H 2 SO 4 sulfuric acid 2.7

A base can be defined as a substance that yields hydroxide ions (OH - ) when dissolved in water. NaOH sodium hydroxide KOH potassium hydroxide Ba(OH) 2 barium hydroxide 2.7