PowerPoint Lectures to accompany Physical Science, 6e Chapter 8 Atoms and Periodic Properties Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display. The origins of quantum physics Atomic structure discovered Ancient Greeks • Democritus (460-362 BC) - indivisible particles called “atoms” • Prevailing argument (Plato and Aristotle) matter is continuously and infinitely divisible John Dalton (early 1800’s) - reintroduced atomic theory to explain chemical reactions Dalton’s atomic theory 1. All matter = indivisible atoms 2. An element is made up of identical atoms 3. Different elements have atoms with different masses 4. Chemical compounds are made of atoms in specific integer ratios 5. Atoms are neither created nor destroyed in chemical reactions Discovery of the electron J. J. Thomson (late 1800’s) • Performed cathode ray experiments • Discovered negatively charged electron • Measured electron’s charge-to-mass ratio • Identified electron as a fundamental particle • What appears to be visible light coming through the slit in this vacuum tube is produced by cathode ray particles striking a detecting screen. You know it is not light, however, since the beam can be pulled or pushed away by a magnet and since it is attracted to a positively charged metal plate. These are not the properties of light, so cathode rays must be something other than light. Electron charge and mass Robert Millikan (~1906) • Studied charged oil droplets in an electric field • Charge on droplets = multiples of electron charge • Charge + Thomson’s result gave electron mass Early models of the atom • Dalton - atoms indivisible • Thomson and Millikan experiments – Electron mass very small, no measurable volume – What is the nature of an atom’s positive charge? • Thomson’s “Plum pudding” model – Electrons embedded in blob of positively charged matter like “raisins in plum pudding” The nucleus Ernest Rutherford (1907) • Scattered alpha particles off gold foil • Most passed through without significant deflection • A few scattered at large angles • Conclusion: an atom’s positive charge resides in a small, massive nucleus • Later: positive charges = protons • James Chadwick (1932): also neutral neutrons in the nucleus • From measurements of alpha particle scattering, Rutherford estimated the radius of an atom to be 100,000 times greater than the radius of the nucleus. This ratio is comparable to that of the (A) thickness of a dime to the (B) length of football field. – In 1917 Rutherford broke up the nucleus of the nitrogen atom by bombarding it with alpha particles and was able to identify a particle with a positive charge called a proton. • He also thought that there were neutral particles in the nucleus called neutrons. They were later identified by James Chadwick • The atom has a tiny, massive nucleus made up of protons and neutrons. • Negatively charged electrons, whose charge balances the charge on the protons, move around the nucleus at a distance of about 100,000 times the radius of the nucleus. • atomic number is the number of protons in the nucleus. All atoms of the same element have the same atomic number. Every element has a distinctive atomic number that identifies it The nuclear atom • Atomic number – Number of protons in nucleus – Elements distinguished by atomic number – 113 elements identified – Number of protons = number of electrons in neutral atoms • Isotopes – Same number of protons; different number of neutrons Atomic symbols and masses Mass number • Number of protons + neutrons Atomic mass units (u) • 1/12 of carbon-12 isotope mass Atomic weight • Atomic mass of an element, averaged over naturally occurring isotopes Classical “atoms” Predictions of classical theory • • • • • Electrons orbit the nucleus Curved path = acceleration Accelerated charges radiate Electrons lose energy and spiral into nucleus Atoms cannot exist! Experiment - atoms do exist New theory needed The quantum concept • Max Planck (1900) – Introduced quantized energy • Einstein (1905) – Light made up of quantized photons • Higher frequency photons = more energetic photons Atomic spectra Blackbody radiation • Continuous radiation distribution • Depends on temperature of radiating object • Characteristic of solids, liquids and dense gases Line spectrum • Emission at characteristic frequencies • Diffuse matter: incandescent gases • Illustration: Balmer series of hydrogen lines Bohr’s theory Three rules: 1. Electrons only exist in certain allowed orbits 2. Within an orbit, the electron does not radiate 3. Radiation is emitted or absorbed when changing orbits Quantum theory of the atom • Lowest energy state = “ground state” • Higher states = “excited states” • Photon energy equals difference in state energies • Hydrogen atom example – Energy levels – Line spectra • These fluorescent lights emit light as electrons of mercury atoms inside the tube gain energy from the electric current. As soon as they can, the electrons drop back to their lower-energy orbit, emitting photons with ultraviolet frequencies. Ultraviolet radiation strikes the fluorescent chemical coating inside the tube, stimulating the emission of visible light. Quantum mechanics • Bohr theory only modeled the line spectrum of H • Further experiments established waveparticle duality of light and matter – Young’s two slit experiment produced interference patterns for both photons and electrons Matter waves Louis de Broglie (1923) • Postulated matter waves • Wavelength related to momentum • Matter waves in atoms are standing waves Wave mechanics • Developed by Erwin Schrodinger • Treats atoms as three dimensional systems of waves • Contains successful ideas of Bohr model and much more • Describes hydrogen atom and many electron atoms • Forms our fundamental understanding of chemistry • Wave Mechanics – Electrons do emit light in certain wavelengths based on their energy levels (orbital radius) – Since waves spread out from the electron, the wave mechanic model predicts an area where an electron would be found, and not a specific place where it would be found. The quantum mechanics model • Quantum numbers specify electronic quantum states • Visualization - wave functions and probability distributions • Electrons delocalized Electronic quantum numbers in atoms 1. Principle quantum number, n – Energy level – Average distance from nucleus 2. Angular momentum quantum number, l – Spatial distribution – Labeled s, p, d, f, g, h, … 3. Magnetic quantum number – Spatial orientation of orbit 4. Spin quantum number – Electron spin orientation • The Quantum Mechanics Model – Quantum mechanics describes the energy levels of an electron wave with four quantum numbers. • • • • distance from nucleus (n) energy sublevel (l) orientation in space. (m) direction of spin (s) – Principal quantum number (n) • Describes main energy level of the electron in terms of its distance from the nucleus. • n = 1, 2, 3, 4, 5, 6, 7 – Angular momentum quantum number ( l ) • Defines energy sublevels within the main energy levels • s, p, d, or f designating the type of orbital and also the orbital shape. • The Heisenberg Uncertainty Principle states that you cannot measure the momentum and exact position of an electron at the same time. –What you can measure is the probability that an electron will be found in a certain area, called an orbital. • (A)An electron distribution sketch representing probability regions where an electron is most likely to be found. (B) A boundary surface, or contour, that encloses about 90 percent of the electron distribution shown in (A). This three-dimensional space around the nucleus, where there is the greatest probability of finding an electron, is called an orbital. – Magnetic quantum number ( m) • Defines the orientation in space of the orbitals relative to a magnetic field. • The s orbital has one orientation • The p sublevel can have 3 orientations • The d sublevel can have 5 orientations • The f sublevel can have 7 orientations. • (A) A contour representation of an s orbital. (B) A contour representation of a p orbital. – Spin quantum number ( s ) • Describes the direction of spin of an electron in its orbit. • Electrons occur in pairs and each of the orientations for a sublevel can have one electron pair. • Experimental evidence supports the concept that electrons can be considered to spin one way or the other as they move about an orbital under an external magnetic field. – Pauli Exclusion Principle • No two electrons can have the same set of quantum numbers. • At least one of the quantum numbers must differ. Electron configuration • Arrangement of electrons into atomic orbitals • Principle, angular momentum and magnetic quantum numbers specify an orbital • Specifies atom’s quantum state • Pauli exclusion principle – Each electron has unique quantum numbers – Maximum of two electrons per orbital (electron spin up/down) • Chemical properties determined by electronic structure Writing electron configurations • Electrons fill available orbitals in order of increasing energy • Shell capacities – – – – s=2 p=6 d = 10 f = 14 • Example: strontium (38 electrons) 1s 22s 22p63s 23p64s 23d 104p65s 2 • Electron Configuration – This is a shorthand designation for electron orientation. – The lowest possible energy level is n=1. • If one electron already occupies this energy level, a second can only occupy it if it has a different spin quantum number. – Electron configurations tells you the quantum numbers of the electron. – Energy sublevel Principle quantum number 1s2 two electrons • There are three possible orientations of the p orbital, and these are called px, py, and pz. Each orbital can hold two electrons, so a total of six electrons are possible in the three orientations; thus the notation p6. Periodic chemical properties • Understood in terms of electron configuration • Electrons in outer orbits determine chemical properties • Summarized in the periodic table • Rows = periods • Columns = families or groups – – – – Alkali metals (IA) Alkaline earths (IIA) Halogens (VIIA) Noble gases (VIIIA) • A-group families = main group or representative elements • B-group = transition elements or metals The periodic table Metals, nonmetals and semiconductors • Noble gases - filled shells, inert • 1-2-3 outer electrons – Lose to become positive ions – Metals • 5-7 outer electrons – Tend to gain electrons and form negative ions – Nonmetals • Semiconductors intermediate between metals and nonmetals •Elements and the Periodic Table • Classification is arranging items into groups or categories according to some criteria. • Classifying helps organize your thinking & reveals patterns that might go unnoticed. • The act of classifying creates a pattern that helps you recognize and understand the behavior of fish, chemicals, or any matter in your surroundings. • These fish, for example, are classified as salmon because they live in the northern Pacific Ocean, have pinkish colored flesh, and characteristicall y swim from salt to fresh water to spawn. • Classifying Matter • Matter is usually defined as anything that has mass and occupies space. • Metals and Nonmetals – A metal had the following properties. • • • • • Metallic luster High heat and electrical conductivity. Malleability, able to be rolled or pounded into a thin sheet. Ductile, can be pulled into a wire. all metals are solids at room temp. except mercury, which is is a liquid • Most matter can be classified as metals or nonmetals according to physical properties. Aluminum, for example, is a lightweight kind of matter that can be melted and rolled into a thin sheet or pulled into a wire. Here you see aluminum pop cans that have been compressed into1,600 lb bales for recycling, destined to again be formed into new pop cans, aluminum foil, or perhaps aluminum wire. – A nonmetal has the following properties • No metallic luster • Poor conductor of heat and electricity. • When it is a solid it is brittle so it cannot be pounded or pulled into a wire. • Occur as solids, liquids, or gases at room temperature *Chemistry is the science concerned with the study of the composition, structure, and properties of substances and the transformations they undergo. • Solids, Liquids, and Gases – Gases have no defined shape or defined volume • Low density – Liquids flow and can be poured from one container to another • Indefinite shape and takes on the shape of the container. – Solids have a definite volume • Have a definite shape. • (A)A gas dispenses throughout a container, taking the shape and volume of the container. (B) A liquid takes the shape of the container but retains its own volume. (C) A solid retains its own shape and volume. • Mixtures and Pure Substances – A mixture has unlike parts and a composition that varies from sample to sample. Can be separated into parts by physical means, involving physical changes. Can be separated into pure substances. – A heterogeneous mixture has physically distinct parts with different properties. – A homogeneous mixture is the same throughout the sample – Pure substances are substances with a fixed composition • A classification scheme for matter. – A physical change is a change that does not alter the identity of the matter. – A chemical change is a change that does alter the identity of the matter. – A compound is a pure substance that can be decomposed by a chemical change into simpler substances with a fixed mass ratio (There are millions of different compounds) – An element is a pure substance which cannot be broken down into anything simpler by either physical or chemical means. (There are 115 known elements) • Sugar (A) is a compound that can be easily decomposed to simpler substances by heating. (B) One of the simpler substances is the black element carbon, which cannot be further decomposed by chemical or physical means. • Elements • Names of Elements – The first 103 elements have internationally accepted names, which are derived from: • The compound or substance in which the element was discovered • An unusual or identifying property of the element • Places, cities, and countries • Famous scientists • Greek mythology • Astronomical objects. • Here are some of the symbols Dalton used for atoms of elements and molecules of compounds. He probably used a circle for each because, like the ancient Greeks, he thought of atoms as tiny, round hard indivisible spheres. – Chemical Symbols • There are about a dozen common elements that have a single capitalized letter for their symbol • The rest, that have permanent names have two letters. – the first is capitalized and the second is lower case – the lower case letter is either the second letter in the name, or the letter of a strong consonant heard when the name of the element is spoken • Some elements have symbols from their Latin names. • Ten of the elements have symbols from their Latin names and one element with a symbol from a German name. • a symbol both identifies a specific element and represents an atom of that element • The elements of aluminum, Iron, Oxygen, and Silicon make up about 88 percent of the earth's solid surface. Water on the surface and in the air as clouds and fog is made up of hydrogen and oxygen. The air is 99 percent nitrogen and oxygen. Hydrogen, oxygen, and carbon make up 97 percent of a person. Thus almost everything you see in this picture us made up of just six elements. – Symbols and Atomic Structure • A molecule is a particle that is composed of two or more atoms held together by a chemical bond. • Isotopes are atoms of an element with identical chemical properties, but different masses due to a difference in the number of neutrons. (Frederic Soddy) • The atomic weight of an element is the weighted average of all the masses of the isotopes. –an isotopes contribution is determined by its relative abundance. – Aston provided evidence of isotopes with the mass spectrograph. • A mass spectrum of chlorine from a mass spectrometer. Note that that two separate masses of chlorine atoms are present, and their abundance can be measured from the signal intensity. The greater the signal intensity, the more abundant the isotope. • The atomic mass of an element is the mass of the element compared to an isotope of carbon Carbon 12. – Carbon 12 is assigned an atomic mass of 12.00 u. • The number of protons and neutrons in an atom is its mass number. • Atomic numbers are whole numbers, identify the element (number of protons, and electrons) • Mass numbers are whole numbers, used to identify isotopes.( closest whole number to atomic mass) • The atomic mass of an isotope is not a whole number, with exception of Carbon-12. • Atomic weights are not whole numbers, due to being a weighted average of the masses of the isotopes. • The Periodic Law • Dmitri Mendeleev gave us a functional scheme with which to classify elements. – Mendeleev’s scheme was based on chemical properties of the elements and their atomic weights. – Meyer’s scheme was based on physical properties of the element and their atomic weights. – It was noticed that the chemical properties of elements increased in a periodic manner. – The periodicity of the elements was demonstrated by Medeleev when he used the table to predict the occurrence and chemical properties of elements which had not yet been discovered. Mendeleev is given credit for developing the periodic table, due to these dramatic predictions. • Mendeleev left blank spaces in his table when the properties of the elements above and below did not seem to match. The existence of unknown elements was predicted by Mendeleev on the basis of the blank spaces. When the unknown elements were discovered, it was found that Mendeleev had closely predicted the properties of the elements as well as their discovery. • The Periodic Law – Similar physical and chemical properties recur periodically when the elements are listed in order of increasing atomic number. – the number of electrons around the nucleus is the essential basis for the modern periodic table. This fact was made known after the work of Rutherford And Moseley on the atomic nucleus. • The Modern Periodic Table • Introduction – The periodic table is made up of rows of elements and columns. – The arrangement has a meaning about atomic structure and about chemical behavior. For the table to be meaningful, one must understand the code of this structure. – An element is identified by its chemical symbol. – The number above the symbol is the atomic number – The number below the symbol is the atomic weight of the element. – A row is called a period – A column is called a family • (A) Periods of the periodic table, and (B) families of the periodic table. • Periodic Patterns – The chemical behavior of elements is determined by its electron configuration – Energy levels are quantized so roughly correspond to layers of electrons around the nucleus. (resemble the layers of an onion) – A shell is all the electrons with the same value of n. • n is a row in the periodic table. – The first three periods contain just A families. Each period begins with a single electron in a new outer electron shell. – Each period ends with a completely filled outer shell that has the maximum number of electrons for that shell. – The number identifying the A families identifies the number of electrons in the outer shell, except helium – The outer shell electrons are responsible for chemical reactions. Elements in the same family have the same number of outer shell electrons; so they will have similar chemical properties. – Group A elements are called representative elements – Group B elements are called transition elements. • Chemical Families – IA are called alkali metals because the react with water to from an alkaline solution, very soft metals. (except H2) – Group IIA are called the alkali earth metals because they are reactive, but not as reactive as Group IA. • They are also soft metals, though not as soft as alkali metals – Group VIIA are the halogens • These need only one electron to fill their outer shell • They are very reactive.( disinfectants, bleach, headlights) – Group VIIIA are the noble gases as they have completely filled outer shells • They are almost non reactive. • Four chemical families of the periodic table: the alkali metals (IA), the alkaline earth metals (IIA), halogens (VII), and the noble gases (VIIIA). • Metals, Nonmetals, and Semiconductors – Chemical behavior is determined by the outer electrons. • These are called valence electrons – These outer shell electrons are represented using electron dot diagrams. – The noble gases have completely filled outer shells and are therefore stable. • All other elements react so as to fill their outer shell and become more stable. • Electron dot notation for the representative elements. – When an atom or molecule gain or loses an electron it becomes an ion. • A cation has lost an electron and therefore has a positive charge • An anion has gained an electron and therefore has a negative charge. • (A) Metals lose their outer electrons to acquire a noble gas structure and become positive ions. (B) Nonmetals gain electrons to acquire an outer noble gas structure and become negative ions. – Elements with 5 to 7 electrons in their outer shell tend to gain electrons to fill their outer shell and become anions. • These are the nonmetals which always tend to gain electrons. – Semiconductors (metalloids) occur at the dividing line between metals and nonmetals. (silicon, germanium, – Elements with 1, 2, or 3 electrons in their outer shell tend to lose electrons to fill their outer shell and become cations. • These are the metals which always tend to lose electrons. • about 80% of all the elements are metals arsenic-conduct electric current under certain conditions) Transition elements are all metals. They are located in B Group families. Form horizontal rows of elements with similar properties. Have variable charges. (Cu+1, Cu+2) • The location of metals, nonmetals, and semiconductors in the periodic table. • The periodic table of the elements.