Chemistry - Chp 4 - Atomic Structure - PowerPoint

Chapter 4 “Atomic Structure”

Section 4.1 Defining the Atom OBJECTIVES: Describe Democritus’s ideas about atoms.

Section 4.1 Defining the Atom OBJECTIVES: Explain Dalton’s atomic theory.

Section 4.1 Defining the Atom OBJECTIVES: Identify what instrument is used to observe individual atoms.

Section 4.1 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”) He believed that atoms were indivisible and indestructible His ideas did agree with later scientific theory, but did not explain chemical behavior, and was not based on the scientific method – but just philosophy

Dalton’s Atomic Theory (experiment based!) 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. 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)

Sizing up the Atom Elements are able to be subdivided into smaller and smaller particles – these are the atoms , and they still have properties of that element If you could line up 100,000,000 copper atoms in a single file, they would be approximately 1 cm long Despite their small size , individual atoms are observable with instruments such as scanning tunneling (electron) microscopes

Section 4.2 Structure of the Nuclear Atom OBJECTIVES: Identify three types of subatomic particles.

Section 4.2 Structure of the Nuclear Atom OBJECTIVES: Describe the structure of atoms, according to the Rutherford atomic model.

Section 4.2 Structure of the Nuclear Atom One change to Dalton’s atomic theory is that atoms are divisible into subatomic particles: Electrons Protons Neutrons

Discovery of the Electron In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle: the electron

Modern C athode R ay T ubes Cathode ray tubes pass electricity through a gas that is contained at a very low pressure. Television Computer Monitor

Mass of the Electron 1916 – Robert Millikan determines the mass of the electron: 1/1840 the mass of a hydrogen atom; has one unit of negative charge The oil drop apparatus Mass of the electron is 9.11 x 10 -28 g

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 be positive particles in the atom to balance the negative charge of the electrons Electrons have so little mass that atoms must contain other particles that account for most of the mass

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 the “ neutron ” – a particle with no charge, but a mass nearly equal to a proton

Subatomic Particles Nucleus 1.67 x 10 -24 0 Neutron (n o ) Nucleus 1.67 x 10 -24 +1 Proton (p + ) Electron cloud 9.11 x 10 -28 -1 Electron (e - ) Location Mass (g) Charge Particle

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

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

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 #1 Target #2

The Answers: Target #1 Target #2

Rutherford’s Findings The nucleus is small The nucleus is dense The nucleus is positively charged 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 Rutherford Atomic Model Based on his experimental evidence: The atom is mostly empty space All the positive charge, and almost all the mass is concentrated in a small area in the center. He called this a “ nucleus ” The nucleus is composed of protons and neutrons (they make the nucleus!) The electrons distributed around the nucleus, and occupy most of the volume His model was called a “ nuclear model ”

Section 4.3 Distinguishing Among Atoms OBJECTIVES: Explain what makes elements and isotopes different from each other.

Section 4.3 Distinguishing Among Atoms OBJECTIVES: Calculate the number of neutrons in an atom.

Section 4.3 Distinguishing Among Atoms OBJECTIVES: Calculate the atomic mass of an element.

Section 4.3 Distinguishing Among Atoms OBJECTIVES: Explain why chemists use the periodic table.

Atomic Number Atoms are composed of identical 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 The “ atomic number ” of an element is the number of protons in the nucleus # protons in an atom = # electrons

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

Mass Number Mass number is the number of protons and neutrons in the nucleus of an isotope: Mass # = p + + n 0 8 8 18 18 Arsenic 75 33 75 Phosphorus 15 31 16 15 - 31 42 33 - 10 Oxygen - Mass # e - n 0 p + Nuclide

Complete Symbols Contain the symbol of the element, the mass number and the atomic number. X Mass number Atomic number Subscript -> Superscript ->

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

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

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

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

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, different mass numbers. These are called isotopes .

Isotopes Frederick Soddy (1877-1956) proposed the idea of isotopes in 1912 Isotopes are atoms of the same element having different masses , due to varying numbers of neutrons. Soddy won the Nobel Prize in Chemistry in 1921 for his work with isotopes and radioactive materials.

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

Isotopes are atoms of the same element having different masses , due to varying numbers of neutrons. 2 1 1 Hydrogen-3 (tritium) 1 1 1 Hydrogen-2 (deuterium) 0 1 1 Hydrogen–1 (protium) Nucleus Neutrons Electrons Protons Isotope

Isotopes Elements occur in nature as mixtures of isotopes. Isotopes are atoms of the same element that differ in the number of neutrons.

Atomic Mass How heavy is an atom of oxygen? It depends, because there are different kinds of oxygen atoms. We are more concerned with the average atomic mass. This is based on the abundance (percentage) of each variety of that element in nature. We don’t use grams for this mass because the numbers would be too small.

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

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

Atomic Masses Atomic mass is the average of all the naturally occurring isotopes of that element. Carbon = 12.011 <0.01% 6 protons 8 neutrons 14 C Carbon-14 1.11% 6 protons 7 neutrons 13 C Carbon-13 98.89% 6 protons 6 neutrons 12 C Carbon-12 % in nature Composition of the nucleus Symbol Isotope

- Page 117 Question Solution Answer Knowns and Unknown

Section 4.4 – The Periodic Table: Organizing the Elements OBJECTIVES Describe the origin of the periodic table Identify the position of groups, periods and the transition metals in the periodic table

Development of the Periodic Table Dmitri Mendeleev (1834 – 1907) Listed the elements in columns in order of increasing mass Then he arranged the columns so that the elements with the most similar properties were side by side

Periodic Table – an arrangement of the elements according to similarities in properties

Henry Moseley (1913) Determined atomic number of the atoms of elements

The Modern Periodic Table Period – Horizontal rows of the periodic table – (side to side) There are 7 periods Periodic law – When the elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties

The Modern Periodic Table Group – vertical (up and down) column of elements in the periodic table The elements in any group of the periodic table have similar physical and chemical properties Each group is identified by a number and letter A or B Group A elements are called representative elements because they exhibit a wide range of both physical and chemical properties

Representative Elements These elements can be divided into three broad classes Metals Characteristics High electrical conductivity High luster when clean Ductile (able to be drawn into wire) Malleable (able to be beaten into sheets)

Representative Elements With the exception of hydrogen all the representative elements on the left side of the periodic table are metals Group 1A – alkali metals Group 2A – alkalaine earth metals Most of the remaining elements that are not in Group A are also metals Transition metals Inner transition metals Group B elements

Rare Earth metals – the inner transition metals which appear below the main body of the periodic table

2. Nonmetals Occupy the upper-right corner of the periodic table Non-metals are elements that are generally non lustrous and poor conductors of electricity Many are gases @ room temperature Ex.) O, Cl Others are brittle solids Ex.) sulfur

2. Nonmetals Group 7A = Halogens ex.) Cl, Br Group 8A = Noble Gases “inert gases” - undergo few chemical changes

3. Metalloids border the line between metals and non metals Have properties in between metals and non metals

Conclusion Rather than memorizing more than 100 elements you need to only learn the general behavior and trends within the major groups This gives you a working knowledge of the properties of most elements

Chemistry - Chp 4 - Atomic Structure - PowerPoint

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