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Chapter 5 “Atomic Structure”

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1 Chapter 5 “Atomic Structure”
Chemistry Olympic High School Mr. Daniel Credits: Stephen L. Cotton Charles Page High School

2 Section 5.1 Defining the Atom
OBJECTIVES: Describe the history of the development of ideas about atoms. Explain Dalton’s atomic theory. Describe the size of the atom.

3 Section 5.1 Defining the Atom
The Greek philosopher Democritus (460 B.C. – 370 B.C.) was among the first to suggest the existence of atoms (from the Greek word “atomos” – meaning “unable to be cut”)

4 Antoine Laurent Lavoisier ( ) The Father of Modern Chemistry: discovered oxygen & hydrogen, developed modern thermodynamics, invented the first periodic table

5 Dalton’s Atomic Theory (experiment based)
All elements are composed of tiny indivisible particles called atoms Atoms of the same element are identical. Atoms of any one element are different from those of any other element. John Dalton (1766 – 1844) Atoms of different elements combine in simple whole- number ratios to form chemical compounds In chemical reactions, atoms are combined, separated, or rearranged – but never changed into atoms of another element.

6 Section 5.2 Structure of the Nuclear Atom
OBJECTIVES: Identify three types of subatomic particles and their properties. Describe the structure of atoms, according to the Rutherford Model.

7 Section 5.2 Structure of the Nuclear Atom
One change to Dalton’s atomic theory is that atoms are divisible into subatomic particles: Electrons, protons, and neutrons are examples of these fundamental particles There are many other types of particles, but we will study these three

8 The electron’s charge-to-mass ratio = 1.76  108 C/g
In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle: the electron JJ Thomsonl The electron’s charge-to-mass ratio = 1.76  108 C/g

9 A Standard Television Tube (First common in the 1950’s)
Cathode Ray Tubes A Standard Television Tube (First common in the 1950’s) A CRT from a Radar Scope (WW II) Standard Computer Monitor (often called a “CRT”) Cathode ray tubes pass electricity through a gas that is contained at a very low pressure.

10 Mass of the Electron simulation Robert Millikan Mass of the electron is 9.11 x g The oil drop apparatus 1916 – Robert Millikan determined the mass of the electron: 1/1840 the mass of a hydrogen atom; has one unit of negative charge

11 Conclusions from the Study of the Electron:
Cathode rays have identical properties regardless of the element used to produce them. All elements must contain identically charged electrons. Atoms are neutral, so there must also be positive particles in the atom to balance the negative charge of the electrons Electrons have very little mass, therefore atoms must contain other particles that account for most of the mass

12 Thomson’s Atomic Model
J. J. Thomson Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model.

13 Other Particle Discoveries:
Eugen Goldstein first observed evidence of what is now called the “proton” in particles with a positive charge, and a mass of 1840 times that of an electron. It’s existance was later confirmed by Ernest Rutherford in 1919. 1932 – James Chadwick confirmed the existence of the “neutron” – a particle with no charge, but a mass nearly equal to a proton

14 The problem: In the following pictures, there is a target hidden by a cloud. To figure out the shape of the target, we shot some beams into the cloud and recorded where the beams came out. Can you figure out the shape of the target? Target #1 The Answer:

15 The Answer Target #2

16 Ernest Rutherford’s Gold Foil Experiment - 1911
Alpha particles are helium nuclei - The alpha particles were fired at a thin sheet of gold foil Particle that hit on the detecting screen (film) are recorded

17 Rutherford’s Findings
Most of the particles passed right through A few particles were deflected VERY FEW were greatly deflected “Like howitzer shells bouncing off of tissue paper!” Conclusions: The nucleus is small The nucleus is dense The nucleus is positively charged

18 The Rutherford Atomic Model
Based on his experimental evidence, Rutherford’s Nuclear Model was developed and stated: The atom is mostly empty space The electrons are distributed around the nucleus, and occupy most of the volume Because of the exceptionally high mass of the nucleus, it must contain particles in addition to protons (neutrons were discovered later) All the positive charge, and almost all the mass is concentrated in a small area in the center. He called this a “nucleus” Ernest Rutherford video

19 Subatomic Particles Particle Charge Mass (g) Location
Crash Course Chemistry - History of Chemical Concepts Particle Charge Mass (g) Location Electron x Electron (e-) cloud Proton x Nucleus (p+) (1 atomic mass unit) Neutron x Nucleus (no) (1 atomic mass unit)

20 Sizing up the Atom Elements can be subdivided into smaller and smaller pieces until there are only single atoms; The smallest particles of an element that still have the properties of that element. 100,000,000 copper atoms in a single file, would be approximately 1 cm long Despite their incredibly small size, individual atoms have recently become observable with scanning tunneling microscopes The Beginning: Xe on Ni Atomic Kanji “original child” Forming a Image: Fe atoms on a Cu surface Image: Fe atoms on a Cu surface Image: Fe atoms on a Cu surface An Artistic View of Ni Carbon Monoxide Man

21 Section 5.3 Distinguishing Among Atoms
OBJECTIVES: Explain what makes isotopes different from each other. Calculate the number of protons, neutrons and electrons in an atom using atomic number and mass Calculate the average atomic mass and atomic number.

22 Atomic Number Element # of protons # of electrons Atomic # (Z) Carbon
Atoms are composed of protons, neutrons, and electrons How then are atoms of one element different from another element? Elements are different because they contain different numbers of PROTONS Atomic number (Z) : the number of protons in the nucleus of each atom of that element. An atom is always neutral, therefore: # protons in an atom = # electrons Element # of protons # of electrons Atomic # (Z) Carbon 6 Phosphorus 15 Iodine 53 Gold 79

23 Mass Number Nuclide p+ n0 e- Mass #
Mass number is the number of protons and neutrons in the nucleus of an atom: Mass # = p+ + n0 Nuclide p+ n0 e- Mass # Oxygen - 10 - 33 42 - 31 15 18 8 8 18 Arsenic 75 33 75 Phosphorus 16 15 31

24 X Complete Symbols Mass U
Complete Symbols contain the symbol of the element, the mass number and the atomic number. Examples: C U Mass number Superscript → 12 X 6 14 6 235 Atomic number Subscript → 92

25 Symbols continued… We can also put the mass number after the name of the element: carbon-12 carbon-14 uranium-235

26 Br 80 35 Symbols Find each of these: number of protons
number of neutrons number of electrons Atomic number Mass Number 35 b) 45 c) 35 d) 35 e) 80

27 Symbols If an element has an atomic number of 34 and a mass number of 78, what is the: number of protons number of neutrons number of electrons complete symbol 34 b) 44 c) 34 d) Se 78 34

28 Symbols If an element has 91 protons and 140 neutrons what is the
Atomic number Mass number number of electrons complete symbol 91 b) 231 c) 91 d) Pa 231 91

29 Symbols If an element has 78 electrons and 117 neutrons what is the
Atomic number Mass number number of protons complete symbol 78 b) 195 c) 78 d) Pt 195 78

30 Isotopes Dalton was wrong about all elements of the same type being identical: Atoms of the same element can have different numbers of neutrons. Thus, they have different mass numbers. These are called isotopes.

31 Isotopes are atoms of the same element having different masses, due to varying numbers of neutrons.
Protons Electrons Neutrons Nucleus Hydrogen–1 (protium) 1 Hydrogen-2 (deuterium) Hydrogen-3 (tritium) 2

32 Isotopes Elements occur in nature as mixtures of isotopes.

33 Atomic Mass How heavy is an atom of oxygen?
It depends, because there are different kinds of oxygen atoms. We generally refer to the average atomic mass. Average atomic mass is based on the abundance (percentage) of each isotope of an element as it is found in nature. It is the number (red) that we find on the periodic table

34 Measuring Atomic Mass Instead of grams, the mass unit we use for atoms is the Atomic Mass Unit (amu or µ) It is defined as one-twelfth the mass of a carbon-12 atom. Carbon-12 is chosen because of its isotope purity. Each isotope has its own atomic mass, thus we determine the average mass from the percent abundance of each isotope.

35 Composition of the nucleus
Atomic Masses Atomic mass is the average of all the naturally occurring isotopes of that element. Isotope Symbol Composition of the nucleus % in nature Carbon-12 12C 6 protons 6 neutrons 98.89% Carbon-13 13C 7 neutrons 1.11% Carbon-14 14C 8 neutrons <0.01% Carbon = µ

36 To calculate the average atomic mass:
Multiply the mass number of each isotope by its abundance (expressed as a decimal), then add the results. The mass of the isotope is expressed in atomic mass units (amu or µ) Isotope Symbol % in nature Carbon-12 12C 98.89% Carbon C 1.11% Carbon C <0.01% 12.01 µ (12.00 µ) (.9889) + (13.00 µ) (.0111) =

37 Atomic Mass Calculation
Element X has two natural isotopes. The isotope with a mass of amu has a relative abundance of 19.91%. The isotope with a mass of amu has a relative abundance of 80.09%. Calculate the atomic mass of this element. ( µ) (.1991) + ( µ) (.8009) = 10.81 µ Boron

38 Section 5.4 Organizing the Elements
OBJECTIVES: Explain how elements are organized in a periodic table. Compare early and modern periodic tables. Identify three broad classes of elements. Distinguish different areas of the periodic table.

39 Section 5.4 Organizing the Elements
A few elements, such as gold and copper, have been known for thousands of years. Yet, only about 13 (of 90) had been identified by the year 1700. As more were discovered, chemists realized they needed a way to organize the elements.

40 Section 5.4 Organizing the Elements
Chemists used the properties of elements to sort them into groups. In 1829 J. W. Dobereiner arranged elements into triads – groups of three elements with similar properties     ( ) ÷ 2 = 88 Ca 40 Sr 88 Ba 137 Li Na 23 K 39 Cl Br I

41 John Newlands ( ) Law of Octaves: noted that after interval of eight elements, similar physical/chemical properties reappeared.  Newlands was the first to formulate the concept of periodicity (repeating patterns) in the properties of the chemical elements.

42 Mendeleev’s Periodic Table
Lothar Meyer Mendeleev’s Periodic Table By the mid-1800s, about 70 elements were known to exist Dmitri Mendeleev – A Russian chemist arranged the elements in order of increasing atomic mass It was the beginning of the modern “Periodic Table”

43 Co and Ni; Ar and K; Te and I
His Completed Work… His Original Table… Co and Ni; Ar and K; Te and I Mendeleev left blanks for elements he predicted that existed such as Germanium But, there were problems: some elements did not fit with their groups When discovered, the elements generally matched his predictions

44 A better arrangement… In 1913, Henry Moseley – British physicist, arranged elements according to increasing atomic number His basic arrangement is still used today The symbol, atomic number & mass are the basic items included in the periodic table

45 Squares in the Periodic Table
The periodic table displays the symbols and names of the elements, along with atomic number and average atomic mass Black symbol = solid Red symbol = Blue symbol = liquid 25º C


47 There are other possible arrangements. How about a:
Spiral Periodic Table

48 The Periodic Table: Your “best friend” is an arrangement of elements in which they are separated into groups based on a set of repeating properties. The periodic table allows you to easily compare the properties of one element to another

49 The Periodic Law Vertical column = group (or family)
Similar physical & chemical properties Identified by number & letter Horizontal rows = periods Repeated properties in each row There are 7 periods When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.

50 Areas of the periodic table
Three classes of elements are: 1) Metals 2) Nonmetals 3) Metalloids

51 Metals Metals: electrical conductors, have luster, are ductile and malleable (in pink below)

52 Nonmetals: generally brittle and dull, are poor conductors of heat and electricity (in blue below)
Some nonmetals… are gases (O, N, Cl); are brittle solids (S, C); one is dark red liquid (Br)

53 Notice the heavy, stair-step line…
Metalloids: border the line Their properties are intermediate between metals and nonmetals (in green below)

54 Elements in the 1A-7A groups are called the Representative Elements
Have very predictable properties and patterns of behavior 1A 8A 2A 3A 4A 5A 6A 7A

55 Groups of elements Group IA – Alkali Metals Alkali Metals
Are the most reactive metals Form a “base” when reacting with water Alkali Metals

56 Groups of Elements Alkaline Earth Metals
Group 2A – Alkaline Earth Metals Are the second most reactive metals Also form bases with water, but do not dissolve well; hence “earth metals” Alkaline Earth Metals

57 Groups of elements The Halogen Group Group 7A – Halogens
The most reactive non-metal group Means “salt-forming” The Halogen Group

58 Groups of elements Noble Gases Group 8A- Noble Gases
Previously called “inert gases” because they rarely take part in a reaction Noble gases have a completely full electron arrangement Noble Gases

59 The “B” groups are called the Transition Elements
La Ac Lanthanide Series Actinide Series The “Inner Transition Metals” actually belong here Crash Course Chemistry - The Periodic Table

60 End of Chapter 5

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