What happens to the size of the metal atoms as you go down group I (alkali metals)?   Why is this so?
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As you go down group I (alkali metals), the size of the metal atoms will be larger. This is because the number of electrons (the atomic number) of each atom gets larger as you go down, thus the electrons occupy more shells, thus they cause the diameter of the atom larger.
846039
The size increases because the electron orbits add up as you go down the group.
495461
Atomic radius (or size) increases as you go down and to the left on the periodic table.  Therefore, out of these, the largest would be K.
961493
You need to know more than just the electron configuration.  There is contraction from left to right on the periodic table, plus additional periodic trends that must be considered.   \n\nEven more importantly, you need to define whether you are looking for the largest element IN ITS NATURAL  STATE (i.e. Cl2, Hg metal, etc) or AS THE ION.   Atomic and ionic radii differ.\n\nMore clarity is needed here.
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When a Potassium atom becomes a Potassium ion, it looses a single electron to become K+.\nThe one electron it looses is located in its outermost, valance, electron shell (the S orbital).\nWhen this electron is taken away, the S orbital is empty and the radius of the resulting ion is now smaller than the previous atom.
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Alkali earths are the elements of group 2.\nAlkali metals are the elements of group 1.\n\nHere are the properties of group 2\nhttp://en.wikipedia.org/wiki/Alkaline_earth_metal\n\nHere are the properties of group 1\nhttp://en.wikipedia.org/wiki/Alkali_metal\n\nYou can compare them.\n\nHere is the water reaction with elements of group 1 (example K)\n\n2K (s) + 2H2O (l) → 2KOH (aq) + H2 (g)\n\nHere is the water reaction with elements of group 2 (example Mg)\n\nMg + 2H2O → Mg(OH)2 + H2\n\nAs you go to the extremes of the table and down the table the reactivity is higher because the electronegativiy is lower so they have a greater tendency to lose the electron and react.
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I'm not directly giving you the answers, you have to know how to do this stuff, but I'll give you the process to get to each.\n\n9: Americium and Curium (wikipedia!)\n\n10: Potassium ion wants to be like the noble gas closest to it (the one right before it), so it loses 1e- to get there and be K+.  If that's not a choice, then look for another ion, positive or negative, that adds up to the electronic configuration of the aforementioned noble gas. (1s2 2s2 2p6 3s2 3p6), 18 total e-'s...\n\n11: Mg is in column 2, things in the same period tend to be chemically similar.  Find the other one in the choices that's in that column\n\n12: U235 and U238 are isotopes, same atomic number, different weight.  If you change the number of protons, you change the atomic number, so that wouldn't be U anymore.  Changing electrons can make ions, but not change weight.  So different number of neutrons for the same element makes an isotope.\n\n13: Same explanation as above, it'll also be Hydrogen's atomic number because it's an isotope.\n\n14: shell 2 contains the 2s and 2p subshells.  2s can hold 2 because it's an s orbital.  2p can hold 2 per p orbital, and a p subshell has 3 p orbitals, so 6.  Total would then be...\n\n15: If an atom loses electrons, it's LOSING -1 per electron lost, so it's becoming more positive.\n\n16: Which element in that list only has 1s shell? Which one has a  stable ion as just a proton or with a configuration like Helium (1s2)?\n\n17: Inert elements are on the far right of the periodic table.  They are inert because they have 8 valence (outer shell) electrons (or 2 in Helium's case) so they can't take anymore and are stable.  Look at your periodic table or count up electrons for each of those atomic numbers.  Only one of them will have 8 in the outer shell...
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For this you should see the electronic configuration of those elements.\nThe metals with atomic number 21-30 are transition metals; they have outer 4s shell complete and the inner 3d orbital is 'in the process' of being filled up. They all have 2 outermost electrons in the 4s2 shell. \nIonization energy is the energy required in removing an outer most electron from a neutral atom. The transition metals and calcium have their outermost 2 electrons in 4s which means that their electrons experience the same force of attraction from the nucleus (being at the same distance from the nucleus) ; hence almost the same amount of energy will be required to overcome that attractive force to remove the electrons.
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The size is decreasing as you go across that list.  A rough rule of thumb is "The more positive the charge, the smaller the ion."   (The reverse is true, too...the more negative the charge, the larger the ion.)
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cations form when the atoms lose an electron or more and have + valency , so the mass will decrease and the effect of the nuclei on the remaining electrons will be stronger so it attracts them stronger toward the center so the size will decrease .. while in the anions the number of electrons increase so the attraction caused by the nuclei is weaker so the size is larger ..\n\nsize of the cation of an element < size of the atom of that element < size of the anion of that element
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a typical metalic atom takes a positive charge when it becomes an ion.  that means that it loses its valence electrons.  that means the radius is smaller in the ion than in the atom.
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The alkali metals, also known as Group 1.  They are lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs), and francium (Fr).  The single valence electron is so easily given away that the elements lower down in the group react explosively with water.  Because they are so highly reactive, they are not found in elemental form in nature, only as ionic compunds.  They are stored under mineral oil to prevent reactions with the moisture in air.\n\nHydrogen is also in Group 1, but its single electron is held very close to the nucleus (a single proton), which means that the electron is not easily given away.  Hydrogen can form a positive or negative ion, because it can also easily accept a second electron to fill its valence shell.  Thus, it is not an alkali metal and is not a member of the family you are asking about.  However, under extremely high pressure, as occurs near the core of Jupiter, hydrogen does take on a metallic structure and behaves as an alkali metal.
68791
Reactivity in goup 1 increses down the group from Lithium to Sodium to Potassium. Therefore, Sodium would the closest. They are all group 1 metals, just in different periods. All have one vacant electron which makes them very reactive.
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They would lose electrons.  metals lose electrons to obtain there octet.  \n\nNa, K would lose 1 electron\n\nMg and Ca would lose 2.\n\n\nwhoever said Ca would gain an electron is an idiot
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The atomic radii for Cr, Mn, Fe, Co, Ni, and Cu are\n\nCr: 125 pm\nMn: 137 pm\nFe: 124 pm\nCo: 125 pm\nNi: 125 pm\nCu: 128 pm\n\nSo, you can see, there is a bit difference between their sizes. Although among them, Cr, Co, Ni, and Cu do have very similar atomic radius. \n\nThe reason that these elements have similar radius is due to their electron configurations. So, let's look at their corresponding electron configuration first (hope you know what I am referring to) :\n\nCr: [Ar] 3d5 4s1\nMn: [Ar] 3d5 4s2\nFe: [Ar] 3d6 4s2\nCo: [Ar] 3d7 4s2\nNi: [Ar] 3d8 4s2\nCu: [Ar] 3d10 4s1\n\nAs you may realize, the major difference between their electron configuration is mainly in the "3d" orbital, which has a quantum number (n) of 3; this 3d orbital has a smaller radius than the next orbital, 4s. The reason why 4s orbital is filled BEFORE 3d is due to the lower energy level of 4s compared to 3d.\n\nIn short, the major reason why these elements have similar radii is because the newly added electrons are not added to the "outermost electron shell", instead, they are added to the "inner electron shell". Therefore, no SIGNIFICANT changes to their radius.
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if u are referring to expansion due to heating than the order of expansion of the substances from lowest to largest is solid, liquids, than gases. But a solid's expansion is extremely low, almost impossible to measure unless its a metal.
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GIVE ME TEN POINTS!!!  \n\nHERES UR GROUPS....\n\n\nGroup 0 Elements\nThe elements in this group have a full outer electronic shell and thus, the elements in this group have no tendency to lose, gain or share electrons. Thus, the elements in the group are chemically inert. \n\nAll the elements in this group are gaseous. Because of their chemical inertness, the elements in this group are called the Nobel Gases : \n\n\nHelium\nNeon\nArgon\nKrypton\nXenon\nRadon \n\n\n--------------------------------------------------------------------------------\n\nGroup I Elements\nThe elements in this group have one electron in their outer electronic shell. Thus, each element in this group has a tendency to lose a single electron, to form a singly charged positive ion, which has the stable electronic configuration of its neighbouring Nobel Gas element in the periodic table. \n\nThe elements in the group are chemically reactive. \n\nThe group is divided into two sub-groups. \n\nSub-Group Ia : The Alkali Metals\n\nLithium\nSodium\nPotassium\nRubidium\nCaesium\nFrancium\n\n\nSub-Group Ib : Transition Metal Elements\n\nCopper\nSilver\nGold\n\nHydrogen is included in this group because it has a single electron in its outer electronic shell. However, hydrogen has none of the metallic properties of the alkali metals. \n\n\n\n--------------------------------------------------------------------------------\n\nGroup II Elements\nThe elements in this group have two electrons in their outer electronic shell. Thus, each element in this group has a tendency to lose two electron, to form a doubly charged positive ion, which has the stable electronic configuration of its neighbouring Nobel Gas element in the periodic table. \n\nThe elements in the group are chemically reactive. \n\n\nSub-Group IIa : The Alkaline Earth Metals\n\nBeryllium\nMagnesium\nCalcium\nStrontium\nBarium\nRadium\n\n\nSub-Group IIb : Transition Metal Elements\n\nZinc\nCadmium\nMercury\n\n\n\n--------------------------------------------------------------------------------\n\nGroup III Elements\nThe elements in this group have three electrons in their outer-most electronic shell. The electronic configurations of the n th orbitals (i.e. outer-most orbitals) are ns2 np1. Thus, each element in this group has a tendency to lose three electron, to form a triply charged positive ion, which has the stable electronic configuration of its nearest neighbouring Nobel Gas element in the periodic table. \n\nThe elements in the group are chemically reactive. \n\n\nSub-Group IIIa : Transition Metal Elements\n\nScandium\nYttrium\nLanthanum\n\nThe following elements also have the same electronic configuration as lanthanum (i.e. 4s2 4p1) in their outer-most electronic orbitals, while the inner 3d orbitals are being filled, from going from element to element.\nCerium\nPraeseodymium\nNeodymium\nPromethium\nSamarium\nEuropium\nGadolinium\nTerbium\nDysprosium\nHolmium\nErbium\nThulium\nYtterbium\nLutecium\n\n\nThese are the transition elements, and are also called the Lanthanides. \n\nActinium\n\nThe following elements also have the same electronic configuration as actinium (i.e. 5s2 5p1) in their outer-most electronic orbitals, while the inner 4d orbitals are being filled, on going from element to element.\nNeptunium\nPlutonium\nAmericium\nCurium\nBerkelium\nThese are the inner transition elements, and are also called the Actinides. \n\n\nSub-Group IIIb : Main Group Elements\n\nBoron\nAluminium\nGallium\nIndium\nThallium\n\n\n\n--------------------------------------------------------------------------------\n\nGroup IV Elements\nThe elements in this group have four electrons in their outer electronic shell. Thus, each element in this group has a tendency to share these four electrons, to form covalent compounds, thereby gaining the stable electronic configuration of its neighbouring Nobel Gas element in the periodic table. \n\nThe elements in the group are chemically reactive. \n\nSub-Group IVa : Transition Me