Electron Configuration and Periodic Trends - Chapter 5 section 3 Guided Notes There are several important atomic characteristics that show predictable that you should know. Atomic Radius The first and most important is atomic. Radius is the from the center of the nucleus to the edge of the cloud. Since a cloud s edge is difficult to define, scientists use the distance between the nuclei of 2 atoms. Atomic radii are usually measured in picometers (pm) or angstroms (Å). An is 1 x 10-10 m. Example: Two Br atoms bonded together are 2.86 angstroms apart. So, the radius of each atom is 1.43 Å. The trend for atomic radius in a column is to go from at the top to at the bottom of the family. With each step the family, we add an entirely new Periodic energy to the electron, making the atoms with each step. The trend across a horizontal is less obvious. What happens to atomic structure as we step from left to right? Each step adds a and an (and 1 or 2 neutrons). are added to existing Periodic energy levels or sublevels. The effect is that the more nucleus has a greater on the cloud. The nucleus is more and the electron cloud is more. The attraction pulls the cloud in, making atoms as we move from to across a period. What keeps electrons from simply flying off into space? Effective charge is the pull that an electron feels from the. The an electron is to the nucleus, the more pull it feels. As effective nuclear charge, the electron cloud is pulled in. Summarize the trend Shielding As more Periodic levels are added to atoms, the inner of electrons shield the electrons from the. The effective nuclear on those outer electrons is, and so the outer electrons are tightly held. Page1
Electrons the nucleus and valence electrons each other making the atom. When going down a group goes. Since electrons are added from the nucleus, there is attraction. This is due to additional levels and the effect. Each energy level shields the from being pulled in the nucleus. When going across a Size goes. Which is Bigger? 1) Na or K? 2) Na or Mg? 3) Al or I? Atomic Radius Summary Period Trend - Gradual decrease in atomic radii from to. Example: Lithium is larger than Fluorine. Why smaller to the right? Increased charge without shielding pulls electrons in. Group Trend - Gradual increase a group. Example: Iodine is larger than fluorine. Why larger going down? Higher energy levels have larger. Shielding - electrons block the between the nucleus and the valence electrons. The more shells the greater the shielding effect. Rank the following elements in order of increasing atomic radius: 1) Li, Cs, Fr, Na 2) P, N, Sb, As 3) Si, P, Cl, Al 4) F, Cl, Br, I Page2
The Octet Rule The goal of most atoms (except H, Li and Be) is to have an or group of 8 in their energy level. They may accomplish this by either electrons away or them. Metals generally electrons, take them from other atoms. Atoms that have gained or lost are called. Ions When an atom gains an, it becomes charged (more electrons than protons ) and is called an. In the same way that nonmetal atoms can electrons, metal atoms can electrons. They become charged. Ionic Radius are always than the atom. The entire Periodic energy level is during ionization. ions-cations. Positive ions are than the they come from. There is nuclear charge acting on the. Conversely, anions are always than the original atom. Electrons are added to the outer energy level. ions-anions. Negative ions are than the atoms they come from. The nuclear force cannot all the. Cation Formation Na atom - 1 valence electron Effective nuclear charge on remaining electrons increases. Result: a smaller sodium cation, Na + Valence e- lost in ion formation Remaining e- are pulled in closer to the nucleus. Ionic size decreases. Page3
Anion Formation (Now you draw this one - add captions at the bottom. Independent Practice P. 152 Questions 1-3 1. 2. 3. Ionic Radii Summary What is it?? The radii of the most common ions of the elements. Positive ions Negative ions Periodic Trend Cation radii from left to right Anion radii from left to right Group Trend Gradual in ionic radii as you go a group. Which is bigger? Cl or Cl -? K + or K? Ca or Ca +2? I - or Br -? Page4
Ionization Energy This is the important periodic trend. If an electron is given enough (in the form of a ) to the effective charge the electron in the, it can the atom. The has been or. The number of protons and electrons is no longer. The required to remove an electron from an atom is energy. (measured in, kj). The the atom is, the easier its are to remove. Ionization energy and atomic radius are proportional. Ionization energy is always, that is is added to the atom to the electron. Trends in Ionization Energy. Ionization Energy across a period because the positive increases. Metals electrons more easily than nonmetals. Nonmetals lose electrons with (they like to electrons). Ionization Energy UP a group because size increases ( Effect) Which has a higher 1 st ionization energy? 1) Mg or Ca? 2) Al or S? 3) Cs or Ba? Removing electrons from positive ions Second ionization energy energy required to additional electrons from ions. Requires much more than first ionization energy. With each additional electron, the effect of nuclear charge on the electrons, making them increasingly difficult to remove from the shell. Ionization Energy is the required to remove an electron from an atom (in the gas phase). Mg (g) + 738 kj ---> Mg + (g) + e- This is called the FIRST ionization energy because we removed only the electron. This is the SECOND. Mg + (g) + 1451 kj ---> Mg 2+ (g) + e- Ionization Energy Summary What is it? The energy required to remove one from the outer shell of an atom. Ion an atom or group of that bears a positive or negative charge Periodic Trend Ionization Energy from left to right across a period. Example: It takes more energy to remove an electron from than from. Page5
Group Trend Recall from top to bottom down groups. Example: Example: It takes less energy (easier) to remove an outer electron from than from. Rank the following in order of increasing ionization energy: 1) Li, Rb, Fr, Na 2) P, N, Sb, As 3) Si, P, Cl, Al 4) F, Cl, Br, I Electron Affinity What does the word affinity mean? Electron affinity is the energy that occurs when an atom an electron (also measured in kj). Where ionization energy is endothermic, electron affinity is exothermic, but not always. Electron affinity is if there is an or partially empty orbital for an electron to occupy. If there are no empty spaces, a new or periodic energy level must be created, making the process. This is true for the alkaline metals and the gases. Electron Affinity Summary What is it?? The in energy when an electron is added to a atom. Most atoms release energy (F and Cl) Some require energy to gain an electron (N) Period Trend from left to right (halogens gain electrons most readily) Group Trend Generally down the group. Review: energy: the required to remove an from a neutral atom. Electron : the change when a atom attracts an to become a negative. Electronegativity Page6
Electronegativity is a measure of the ability of an atom in a molecule to electrons to itself. Concept proposed by Linus in1901-1994. Electronegativity is a measure of an atom s attraction for another atom s. It is an arbitrary scale that ranges from to 4. The units of are Paulings. Generally, are electron givers and have electronegativities. Nonmetals are electron takers and have electronegativities. What about the noble gases?. In a group: Atoms with energy can attract electrons (less.) So, electronegativity UP a group of elements. In a period: More, while the energy levels are the means atoms can better electrons. So, electronegativity RIGHT in a period of elements. Electronegativity Summary What is it?? The ability of an atom in a compound to attract from another atom in the compound. Periodic Trend Group Trend from left to right from top to bottom is the most electronegative atom. Which is more electronegative? 1) F or Cl? 2) Na or K? 3) Sn or I? Overall Reactivity Page7
This ties all the previous trends together in one package. However, we must treat metals and nonmetals separately. The most reactive metals are the since they are the best givers. The most reactive nonmetals are the ones, the best electron takers. Independent Work p. 164 Questions 1-3 1. a. b. c. d. e. 2. a. b. 3. Periodic Trends Packet Page8