Chapter 0 A Very Brief History of Chemistry Chemistry: The Molecular Nature of Matter, 7E Jespersen/Hyslop

Chapter in Context Scope and purpose of the chemical sciences – Chemistry’s big ideas Formation of the elements – supernovas Distribution of substances around the world – elements and the Earth Atomic theory – explanatory and predictive power The structure of the atom – key experiments

The Four Big Ideas Atomic theory – John Dalton, 1813 Described atoms and how they interact with one another 2. Careful laboratory observation Can lead to understanding the atomic world 3. Energy changes and probability Lead to predictions about chemical interactions Geometric shapes of molecules are important Affect properties, reactivity, and function (e.g. DNA, RNA, and proteins)

Supernovas and the Elements The Big Bang: 14 billion years ago Explosion of energy and subatomic particles Extreme temperature, pressure, and density As the Big Bang cooled Initially only quarks exist After 1 second, quarks form protons and neutrons After 3 minutes, nucleosynthesis begins of light nuclei (e.g., helium, lithium) After further cooling, electrons join nuclei to form atoms.

Supernovas and the Elements Universe was 91% hydrogen, 8% helium, 1% other light atoms Uneven distribution of matter resulted in star formation Formation of elements occurred in the stars Small atoms combined, due to high pressure at the center, to create slightly heavier elements New elements concentrated in the stars’ centers Heavier elements then combined into new, even heavier elements Cycle kept repeating

Supernovas and the Elements Iron is the heaviest element created in stars Causes the nuclear reactions to stop and the star to cool and collapse in on itself allows for even heavier elements to form Eventually, the star disintegrates Called a supernova Spews its content into space Remnants rejoin to form new stars Cycle begins again Some of the debris combines to form moons, planets, and asteroids

Distribution of the Elements Earth formed 4.5 billion years ago Result of gravitational forces Earth heated up Iron and nickel melted Migrated to the core Outer core is superheated lava Mantel is superheated rock Crust is the surface 10 miles thick Contains the familiar elements (gold, silicon, carbon, etc.)

Atomic Theory Most significant theoretical model of nature Atoms Tiny submicroscopic particles Make up all chemical substances Make up everything in macroscopic world Smallest particle that has all properties of given element

Atomic Theory Three important ideas Law of Definite Proportions In a given compound, the elements are always combined in the same proportion by mass Always find 1 g H to 8 g O in water Law of Conservation of Mass No detectable gain or loss of mass occurs in chemical reactions. Mass is conserved. A closed vessel with 16 g O and 2 g H will weigh 18 g after water is formed from them Dalton’s atomic theory

Atomic Theory Law of definite proportions and law of conservation of mass Based on laboratory observations of mass and volume Discussed in detail in a later chapter

Dalton’s Atomic Theory John Dalton Developed underlying theory to explain Law of Conservation of Mass Law of Definite Proportions Reasoned that if atoms exist, they have certain properties

Dalton’s Atomic Theory (cont.) Matter consists of tiny particles called atoms Atoms are indestructible In chemical reactions, atoms rearrange but do not break apart In any sample of a pure element, all atoms are identical in mass and other properties Atoms of different elements differ in mass and other properties In a given compound, constituent atoms are always present in same fixed numerical ratio

Proof of Atoms Early 1980’s, use Scanning Tunneling Microscope (STM) Surface can be scanned for topographical information Image for all matter shows spherical regions of matter Proof of atoms STM of palladium

Discovery of Subatomic Particles Late 1800s and early 1900s Cathode ray tube experiments showed that atoms are made up of subatomic particles Discovered negatively charged particles moving from the cathode to the anode Cathode – negative electrode Anode – positive electrode

Discovery of Electron JJ Thomson (1897) Modified cathode ray tube Made quantitative measurements on cathode rays Discovered negatively charged particles Electrons (e –) Determined charge to mass ratio (e/m) of these particles e/m = –1.76 x 108 coulombs/gram

Millikan Oil Drop Experiment Determining charge on Electron Calculated charge on electron e – = –1.60 × 10–19 Coulombs Combined with Thomson’s experiment to get mass of electron m = 9.09 × 10–28 g

Discovery of Atomic Nucleus Rutherford’s Alpha Scattering Experiment Most alpha () rays passed right through gold A few were deflected off at an angle 1 in 8000 bounced back towards alpha ray source Gave us current model of nuclear atom

Discovery of Proton Discovered in 1918 in Ernest Rutherford’s lab Detected using a mass spectrometer Hydrogen had mass 1800 times the electron mass Masses of other gases whole number multiples of mass of hydrogen Proton Smallest positively charged particle

Rutherford’s Nuclear Atom Demonstrated that nucleus: has almost all of mass in atom has all of positive charge is located in very small volume at center of atom Very tiny, extremely dense core of atom Where protons ( ) and neutrons ( ) are located

Discovery of Neutron First postulated by Rutherford and coworkers Estimated number of positive charges on nucleus based on experimental data Nuclear mass based on this number of protons always far short of actual mass About ½ actual mass Therefore, must be another type of particle Has mass about same as proton Electrically neutral Discovered in 1932 by James Chadwick

Your Turn! Which scientist is correctly matched with the discovery for which he is known? Thomson, neutron Rutherford, electron Chadwick, neutron Milliken, nucleus Dalton, charge on the electron

Atomic Structure Electrons ( , or e –) Very low mass Occupy most of atom’s space Balance of attractive and repulsive forces controls atom size Attraction between protons ( ) and electrons ( ) holds electrons around nucleus Repulsion between electrons helps them spread out over volume of atom In neutral atom Number of electrons must equal number of protons Diameter of atom ~10,000 × diameter of nucleus

Properties of Subatomic Particles Three kinds of subatomic particles of principal interest to chemists Particle Mass (g) Electrical Charge Symbol Electron 9.109  10–28 –1 Proton 1.673  10–24 +1 Neutron 1.675  10–24

Atomic Notation Atomic number (Z) Isotopes Number of protons that atom has in nucleus Unique to each type of element Element is substance whose atoms all contain identical number of protons Z = number of protons Isotopes Atoms of same element with different masses Same number of protons ( ) Different number of neutrons ( )

Atomic Notation Isotope Mass number (A) Atomic Symbols A = (number of protons) + (number of neutrons) A = Z + N For charge neutrality, number of electrons and protons must be equal Atomic Symbols Summarize information about subatomic particles Every isotope defined by two numbers Z and A Symbolized by Ex. What is the atomic symbol for helium? He has 2 e–, 2 n and 2 p Z = 2, A = 4

Isotopes Most elements are mixtures of two or more stable isotopes Each isotope has slightly different mass Chemically, isotopes have virtually identical chemical properties Relative proportions of different isotopes are essentially constant Isotopes distinguished by mass number (A): e.g., Three isotopes of hydrogen (H) Four isotopes of iron (Fe)

Example: What is the isotopic symbol for Uranium- 235? Number of protons ( ) = 92 = number of electrons in neutral atom Number of neutrons ( ) = 143 Atomic number (Z ) = 92 Mass number (A) = 92 + 143 = 235 Chemical symbol = U Summary for uranium-235:

Your Turn! An atom of has ___ protons, ___ neutrons, and ___ electrons. 82, 206, 124 124, 206, 124 124, 124, 124 82, 124, 82 82, 124, 124

Your Turn! What is the correct symbol for an element that has 27 protons, 33 neutrons, and 27 electrons.

Learning Check: 78 53 53 46 35 35 Fill in the blanks: 131I 81Br symbol neutrons protons electrons 131I 81Br 36 29 29 78 53 53 46 35 35

Carbon-12 Atomic Mass Scale Need uniform mass scale for atoms Atomic mass units (symbol u) Based on carbon: 1 atom of carbon-12 = 12 u (exactly) 1 u = 1/12 mass 1 atom of carbon-12 (exactly) Why was 12C selected? Common Most abundant isotope of carbon All atomic masses of all other elements ~ whole numbers Lightest element, H, has mass ~1 u

Calculating Atomic Mass Generally, elements are mixtures of isotopes e.g. Hydrogen Isotope Mass % Abundance 1H 1.007825 u 99.985 2H 2.0140 u 0.015 How do we define atomic mass? Average of masses of all stable isotopes of given element How do we calculate average atomic mass? Weighted average Use isotopic abundances and isotopic masses

Learning Check Naturally occurring magnesium is a mixture of 3 isotopes; 78.99% of the atoms are 24Mg (atomic mass, 23.9850 u), 10.00% of 25Mg (atomic mass, 24.9858 u), and 11.01% of 26Mg (atomic mass, 25.9826 u). From these data calculate the average atomic mass of magnesium. 0.7899 x 23.9850 u = 18.946 u 24Mg 0.1000 x 24.9858 u = 2.4986 u 25Mg 0.1101 x 25.9826 u = 2.8607 u 26Mg Total mass of average atom = 24.3053 u rounds up to 24.31 u

Your Turn! A naturally occurring element consists of two isotopes. The data on the isotopes: isotope #1 68.5257 u 60.226% isotope #2 70.9429 u 39.774% Calculate the average atomic mass of this element. 70.943 u 69.487 u 69.526 u 69.981 u 69.734 u 0.60226 × 68.5257 u = 41.270 u 0.39774 × 70.9429 u = 28.217 u 69.487 u

Your Turn! Magnesium (average atomic mass 24.305 u) consists of three isotopes. isotope #1 23.985 u 78.990% isotope #2 24.986 u 10.000% isotope #3 _______ _______ Calculate the abundance and atomic mass of #3. 25.982 u 24.305 u 24.486 u 11.010 u 17.658 u Abundance: 100.00 – 78.990 – 10.000 = 11.010 % 11.010 % 24.305 u = (0.78990 × 23.985 u) + (0.10000 × 24.986 u) + (0.11010 × X u) Solving for X yields 25.982 u 25.982 u