Presentation is loading. Please wait.

Presentation is loading. Please wait.

Chapter 4 “Atomic Structure”

Similar presentations

Presentation on theme: "Chapter 4 “Atomic Structure”"— Presentation transcript:

1 Chapter 4 “Atomic Structure”

2 Section 4.1 Defining the Atom
Greek philosopher Democritus (460 B.C. – 370 B.C.) suggested existence of atoms (Greek word “atomos”) Believed atoms were indivisible and indestructible His ideas agreed with later scientific theory, but didn’t explain chemical behavior - was not based on scientific methods – only philosophy

3 Dalton’s Atomic Theory (experiment based!)
All elements composed of tiny indivisible particles called atoms Atoms of same element identical. Atoms of any one element are different from all other elements. 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.

4 Sizing up the Atom Elements subdivided into smaller particles – called atoms, and they still have properties of that element 1.0 x 108 copper atoms in a single file, they would be approximately 1 cm long individual atoms are observable with instruments such as scanning tunneling (electron) microscopes

5 Section 4.2 Structure of the Nuclear Atom
One change to Dalton’s atomic theory - atoms are divisible into subatomic particles: Electrons, protons, and neutrons

6 Discovery of Electron J.J. Thomson used cathode ray tube to deduce presence of negatively charged particle…….the electron

7 Modern Cathode Ray Tubes
Television Computer Monitor CRT’s pass electricity through gas contained - very low pressure.

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

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

10 Conclusions from the Study of the Electron:
Eugen Goldstein in 1886 observed what is now called the “proton” - particles with a positive charge, and a relative mass of 1 (or 1840 times that of an electron) 1932 – James Chadwick confirmed the existence of “neutron” – particle with no charge, but mass nearly equal to proton

11 Subatomic Particles Particle Charge Mass (g) Location Electron (e-) -1
9.11 x 10-28 Electron cloud Proton (p+) +1 1.67 x 10-24 Nucleus Neutron (no)

12 Thomson’s Atomic Model
J. J. Thomson Believed electrons were like plums embedded in + charged “pudding,” called “plum pudding” model.

13 Ernest Rutherford’s Gold Foil Experiment - 1911
Alpha particles - helium nuclei w/ + charge - The alpha particles were fired at thin sheet of gold foil Particles that hit on the detecting screen (film) were recorded

14 Rutherford’s 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 #2 Target #1

15 The Answers: Target #1 Target #2

16 Rutherford’s Findings
Most particles passed through Few deflected VERY FEW greatly deflected “Like howitzer shells bouncing off of tissue paper!” Conclusions: Small nucleus Dense nucleus + charge nucleus The Atom Song Atoms song - Mark Rosengarten

17 The Rutherford Atomic Model
His experimental evidence: atom mostly empty space All positive charge, almost all mass in small center. “Nucleus” protons and neutrons make nucleus! electrons distributed around nucleus…occupy most volume His model called “nuclear model” Rutherford’s Atom 3:08

18 # protons in atom = # electrons
Section 4.3 Atomic Number All atoms composed of identical protons, neutrons, and electrons How then are atoms of one element different from another element? Elements different b/c they contain different # of PROTONS “atomic number” of element is number of protons in nucleus # protons in atom = # electrons

19 Atomic Number Atomic number (Z) of element is # of protons in nucleus of each atom of that element. Element # of protons Atomic # (Z) Carbon 6 Phosphorus 15 Gold 79

20 Mass Number Mass number is # of protons and neutrons in nucleus of an isotope: 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

21 Complete Symbols Contain symbol of element, mass number & atomic number. Mass number X Superscript → Atomic number Subscript →

22 Br Symbols 80 35 Find each of these: number of protons
number of neutrons number of electrons Atomic number Mass Number 80 Br 35

23 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

24 Symbols If an element has 91 protons and 140 neutrons what is the Atomic number Mass number number of electrons complete symbol

25 Symbols If an element has 78 electrons and 117 neutrons what is the Atomic number Mass number number of protons complete symbol

26 Isotopes Dalton was wrong about elements of same type being identical… Atoms of same element can have different numbers of neutrons. different mass numbers isotopes

27 Isotopes Frederick Soddy (1877-1956) proposed idea of isotopes in 1912
Isotopes - atoms of same element with different masses, b/c varying #s of neutrons Won 1921 Nobel Prize in Chemistry has a small crater named for him on the far side of the Moon.

28 We can also put mass number after name of the element:
Naming Isotopes We can also put mass number after name of the element: carbon-12 carbon-14 uranium-235

29 Hydrogen–1 (protium) 1 Hydrogen-2 (deuterium) Hydrogen-3 (tritium) 2
Isotope Protons Electrons Neutrons Nucleus Hydrogen–1 (protium) 1 Hydrogen-2 (deuterium) Hydrogen-3 (tritium) 2

30 Isotopes Elements occur in nature as mixtures of isotopes.

31 Atomic Mass How heavy is an oxygen atom?
Depends, b/c different kinds of oxygen atoms exist. We’re more concerned with average atomic mass. Based on abundance (%) of each variety of that element in nature. Don’t use grams - numbers tooooo small.

32 Measuring Atomic Mass Atomic Mass Unit (amu)
one-twelfth mass of a carbon-12 atom. Carbon-12 chosen b/c of its isotope purity. Each isotope has own atomic mass we determine average from % abundance.

33 To calculate the average:
Multiply atomic mass of each isotope by abundance (decimal), then add results. If not told otherwise, mass of isotope expressed in atomic mass units (amu)

34 Composition of the nucleus
Atomic Masses Atomic mass is average of all naturally occurring isotopes of that element. Atomic mass (amu) 12 13.00 14.00 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% What is the average atomic mass of Carbon? 12.01

35 - Page 117 Question Knowns and Unknown Solution Answer

36 End of Chapter 4

Download ppt "Chapter 4 “Atomic Structure”"

Similar presentations

Ads by Google